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962 1962 1
CHEMISTRY SEMESTER 1CHEMISTRY SEMESTER 1CHAPTER 3CHAPTER 3
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CHAPTER 3 CHEMICAL BONDING
31 Ionic bonding
32 Covalent bonding
33 Metallic bonding
34 Intermolecular forces Van der Waals forces andhydrogen bonding
Topic
2007 2008 2009 2010 2011 20122013
Sem 12014
Sem 1
P1 P2 P1 P2 P1 P2 P1 P2 P1 P2 P1 P2 AB
CA
B
C
3
Chemical
Bonding
43c
6a4 5a 3
3b
c1 2
2
5b2
1b
5c
7a
319
b3 19
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INTERACTION
BETWEEN ELEMENTS
Metal and
non-metalNon - metal and
non-metal
Metal and
metal
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bull A Lewis dot symbol consists of the symbol of an element and
one dot for each valence electron in an atom of the element
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31 Ionic Bonding
bull The central idea of the ionic bonding model is the transfer of
electrons from metal atoms to non-metal atoms to form
ions that come together in a solid ionic compound where
ionic bond is formed in between oppositely charged ions
by electrostatic attraction forcesbull For example in the formation of sodium fluoride NaF
Sodium atom Fluorine atom Sodium fluoride
Electronicconfiguration
Na (1s22s22p63s1) F (1s22s22p5) Na+ F-
(1s22s22p6) (1s22s22p6)
Orbitaldiagrams Na F Na+ F-
Lewis
diagram
1s 2s 2p 3s
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bull The interaction between sodium atom and fluorine atom
occur where sodium atom (with low ionisation energy)
donates electron to fluorine (with high electron affinity) to formsodium ion Na+ and fluoride ion F- respectively Note that
both ions have achieved octet arrangement of ns2np6 as it is
the most stable form of an ion formedbull The oppositely charged Na+ and F- form a giant ionic crystal
lattice with very high melting point via electrostatic attraction
i Magnesium chloride MgCl2 ii potassium oxide K2O
iii Calcium sulphide CaS iv aluminium oxide Al2O3
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bull According to Coulumbs Law electrostatic energy between
two oppositely charged substance (A and B) is directly
proportional to the charge carried by each ions yet inversely
proportional to the distance between them
bull This relationship helps us predict trends in lattice energy and
explain the effects of ionic size and charge
( ) minus+
minus+
minus
times
prop r r
energylatticeor EnergyticElectrosta
nn
a) Effect of ionic size As we move down a group in theperiodic table the ionic radius increases Therefore the
electrostatic energy between cations and anions decreases
because the inter-ionic distance is greater thus the lattice
energies of their compounds should decrease as well This
prediction is borne out by the alkali-metal halides note the
regular decrease in lattice energy down a group whether we
hold the cation constant (LiF to LiI)
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b) Effect of ionic charge When we compare lithium fluoride
with magnesium oxide we find cations of about equal radii
(Li+ = 76 pm and Mg2+ = 72 pm) and anions of about equalradii (F- = 133 pm and O2- = 140 pm) Thus the only
significant difference is the ionic charge LiF contains the
singly charged Li+ and F- ions whereas MgO contains the
doubly charged Mg2+ and O2- ions The difference in their
lattice energies is
∆H of LiF = - 1050 kJ mol-1
∆Hlattice of MgO = - 3923 kJ mol-1
This nearly fourfold increase in ∆Hlattice reflects the fourfold
increase in the product of the charges (1 x 1 vs 2 x 2) in the
numerator of equation above
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3 Properties of Ionic Compound
a) Melting point - Ionic compound has giant ionic crystal lattice
which are hold by strong electrostatic attraction forces by
repeating of oppositely charged ions
Therefore the melting point of ionic compounds are usuallyvery high as a lot of energies are required to overcome the
electrostatic attraction forces and melted to form free moving
ions Therefore ionic compounds have very high melting
point
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b) Conductivity of electricity - Most ionic compounds do not
conduct electricity (insulator) in the solid state but do
conduct it when melted or when dissolved in water (exceptsome super-ionic conductors such as AgI which have
remarkable conductivity in the solid state) Solid ionic salt
consists of immobilized ions When it melts or dissolveshowever the ions are free to move and carry an electric
current
-
are hard (does not dent) rigid (does not bend) and brittle(cracks without deforming) These properties are due to the
powerful attractive forces that hold the ions in specific
positions throughout the crystal Moving the ions out of
position requires overcoming these forces so the sample
resists denting and bending If enough pressure is applied
ions of like charge are brought next to each other and
repulsive forces crack the ionic solid suddenly
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32 Covalent bondbull Studies of covalent bond was widely developed ever since
Lewis suggested that a chemical bond exist in a hydrogen gas
occur by sharing en electron between two hydrogen atoms
bull Electron pair that connect the 2 hydrogen atoms is called
covalent bond a bond in which two electrons are shared by
wo a oms an e e ec ron pa r a on e ween e wohydrogen atoms is also called as bonding pair electrons
bull In a covalent bond each electron in a shared pair is attracted to
the nuclei of both atoms This attraction holds the two atoms in
H2 together and is responsible for the formation of covalentbonds in other molecules
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2 A Lewis structure is a representation of covalent bonding
in which shared electron pairs are shown either as lines
or as pairs of dots between two atoms and lone pairsare shown as pairs of dots on individual atoms Only
valence electrons are shown in a Lewis structure
a) Consider the fluorine molecule F2 The electronconfiguration of F is 1s22s22p5 The 1s electrons are in
inner shell which is nearest to the nucleus For this reason
the do not artici ate in bond formation Thus each F
atom has seven valence electrons (2s22p5) Thereforeeach fluorine atom needed one electron to achieve octet
configuration (ns2np6)
F 1 s22s22p5 F 1 s22s22p5 1s2 2s2 2p6 1s2 2s2 2p6
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bull Oxygen atom has electronic configuration of 1s22s22p4 To
achieve stable octet configuration (ns2np6) each oxygen
atom need 2 electrons Hence when 2 oxygen atomsinteract they shared two electrons in between each other as
described in diagram below
O 1s2
2s2
2p4
O 1s2
2s2
2p4
1s2
2s2
2p6
1s2
2s2
2p6
bull From the structure of oxygen molecule formed each oxygen
atom shared two electrons from each other to form a double
bond in order to achieve octet configuration among each
oxygen atom
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bull Nitrogen has electronic configuration of 1s22s22p3 and
required 3 electrons to achieve octet configuration (ns2np6)
In this case each nitrogen atom shared 3 electrons fromeach of its atom to form triple bond
N 1 s22s22p3 N 1 s22s22p3 1s2 2s2 2p6 1s2 2s2 2p6
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Water H2O
H 1 s1 O 2 s22p4
Carbon dioxide CO2
C 2 s22p2 O 2 s22p4
Ammonia NH3
H 1 s1 N 2 s22p3Ethene C2H4
H 1 s1 C 2 s22p2
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Hydrogen cyanide HCN
H 1 s1 N 2 s22p3 C 2 s22p2
Ethanoic acid CH3COOH
H 1 s1 O 2 s22p4 C 2 s22p2
Tetrachloromethane CCl4C 2 s22p2 Cl 3s23p5
Ethyne C2H2
H 1 s1 C 2 s22p2
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a) Note that in ethene hydrogen cyanide and ethanoic acid all
the valence electrons are used in bonding there are no lone
pairs on the carbon atoms In fact most of the stablemolecules containing carbon do not have lone pairs on the
carbon atoms
b) Multiple bonds (double bond or triple bond) are shorter than single covalent bonds Bond length is defined as the
distance between the nuclei of two covalently bonded atoms
nitrogen triple bonds are shorter than double bonds whichin turn are shorter than single bonds The shorter multiple
bonds are also more stable than single bonds
c) Covalent bond not only exist in neutral molecule but also insome molecular ions Table below shows a few example of
molecular ions which have covalent bonds in its molecule
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Carbonate ion CO32-
C 2 s22p2 O 2 s22p4
Cyanide ion CN-
N 2 s22p3 C 2 s22p2
u p a e on 4-
S 3 s23p4 O 2 s22p4
ra e on 3-
N 2 s22p3 O 2 s22p4
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321 Exception of Octet Rules
1 From the Lewis structure sketched for sulphate ion SO42- we can
see that the center metal atom (sulphur) has more than 8electrons For this the molecule is described as molecule that can
expanded octet These molecules that have more than 8
electrons are located at Period 3 and below as these center
atoms have empty d-orbital available to expand the number ofelectrons positioned in the center atom Those center atom from
Period 2 such as C N O and F can only allocate a maximum
Example Both phosphorous and nitrogen are elements from Group15 with the valence electron of ns2np3 They can both form NCl3and PCl3 respectively when react with limited amount of chlorine
however under excess chlorine only PCl 5 can be formed but
not NCl 5 Explain the statement bolded
Solution This is due to phosphorous which is from period
3 have empty d-orbital to expand the octet However
nitrogen which is from Period 2 do not have emptyorbital and can only allocate 8 electrons in its shell
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Phosphorous pentachloride PCl5P 3 s23p3 Cl 3s23p5
Sulphur hexafluoride SF6
S 3 s23p4 F 2 s22p5
Bromine pentachloride BrCl5Cl 3s23p5 Br 4s24p5
Xenon tetrafluoride XeF4
Xe 5s25p6 F 2 s22p5
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2 There are also some stable covalent compounds which have
less than 8 electrons in their center atom (incomplete octet)
The center atoms are usually metals with great number ofvalence electrons with small atomic radius such as
beryllium boron and aluminium
BeCl2Be 2s2 Cl 3s23p5 BF3B 2 s22p1 F 2 s22p5 AlCl3Al 3s23p1 F 3 s23p5
These compound possessed stability due to their short bondlength between center atom with surround atom Furthermore
they can form resonance structure between the center atom
and surrounding atoms
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3 There are also some molecules which contain an odd
number of electrons Among the most common compounds
are nitrogen monoxide (NO) and nitrogen dioxide (NO2)
Odd-electron molecules are sometimes called radicals
Nitrogen monoxide NO
N 2 s22p3 O 2 s22p4
Nitrogen dioxide NO2
N 2 s22p3 O 2 s22p4
any ra ca s are g y reac ve e reason s t at t ere
is a tendency for the unpaired electron to form a
covalent bond with an unpaired electron on another
molecule For example when two nitrogen dioxide
molecules collide they form dinitrogen tetroxide
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322 Dative bond
1 As shown in the Lewis structure of nitrate ion one of the N-Obond is drawn as rarr The bond rarr placed is called as dative
bond (also known as coordinative bond) where dative
bond is defined as a covalent bond in which one of the atoms
donate the lone pair electrons available Although theproperties of a coordinate covalent bond do not differ from
those of a normal covalent bond because all electrons are
a e no ma er w a e r source
2 Dative bond is usually applied for these few circumstances
below
a) To assist atom molecule ion that not yet achieved octet
configuration Making use of atom which has lone pair
electrons to those which are lack of electrons
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Sulphur dioxide SO2Sulphur trioxide SO3 Carbon monoxide CO
Ozone molecule O3 Water with hydrogen ion
S
OO
O
Ammonia with hydrogen ion Ammonia with boron trifluoride
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b) Formation of dimer - In order for some compounds which
have incomplete octet to achieve stability they tend to form
dimer or polymer among themselves by using dative bondTwo of the most common examples are aluminium trichloride
and beryllium dichloride
Monomer of AlCl3Dimer of aluminium chloride Al
2
Cl6
Monomer of BeCl2 Dimer of Be2Cl4 Polymer of (BeCl2)n
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c) Formation of coordination compounds - Coordination
compounds are substances that contain at least one complex
ion a species consisting of a central metal cation (either atransition metal or a main-group metal) that is bonded to
molecules andor anions called ligands via dative
(coordinative) bond
hexaaquacopper (II) ion [Cu(H2O)6]2+ tetraamminenickel (II) ion [Ni(NH3)4]
2+
Hexacyanoferrate (III) ion [Fe(CN)6]3- Trioxalatocobaltate(III)ion [Co(C2O4)3]
3-
3 2 3 Hybridisation Theory
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323 Hybridisation Theory
Valence bond theory
bull The basic principle of valence bond theory is that a covalentbond is formed when orbitals of two atoms overlap and the
overlapped region which is between the nuclei is occupied by
a pair of electrons The central themes of valence bond theory
derive from this principle
ndash Opposing spins of the electron pair ~ Stated in Paulis
orbitals has a maximum capacity of two electrons that musthave opposite spins For example when a covalent bond is
formed in molecule of hydrogen H2 the two 1s electrons of
two H atoms occupy the overlapping 1s orbitals and have
opposite spins
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bull Maximum overlap of bonding orbitals ~ The bond strength
depends on attraction of the nuclei for the shared electrons
so the greater the orbital overlap the stronger (morestable) the bond The extent of overlap depends on the
shapes and directions of the orbitals An s orbital is spherical
but p and d orbitals have more electron density in one
direction than in another Thus whenever possible a bond
involving p or d orbitals will be oriented in the direction that
maximizes overlap For example in hydrogen fluoride (HF)
bond the 1s orbital of H overlaps the half-filled 2p orbital of Falong the long axis of that orbital Any other direction would
result in less overlap and thus a weaker bond
H b idi ti f t i bit l T t f th
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bull Hybridisation of atomic orbitals ~ To account for the
bonding in simple diatomic molecules like HF we picture the
direct overlap of s and p orbitals of isolated atoms But howcan we account for the shapes of so many molecules and
polyatomic ions through the overlap of spherical s orbitals
dumbbell-shaped p orbitals and cloverleaf-shaped d orbitals
Linus Pauling proposed that the valence atomic orbitals in the
molecule are different from those in the isolated atoms The
spatial orientations of these new orbitals lead to more stable
bonds and are consistent with observed molecular shapesThe process of orbital mixing is called hybridisation and
the new atomic orbitals are called hybrid orbitals Two key
points about the number and type of hybrid orbitals are that
i The number of hybrid orbitals obtained equals the number ofatomic orbitals mixed
ii The type of hybrid orbitals obtained varies with the types of
atomic orbitals mixed
3 2 3 1 T f h b idi ti
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3231 Type of hybridisation
1 sp3 hybridisation ~ When four electron groups surround
the central atom the center atom involved must prepare fourorbitals with equal energies to overlap with the four
surrounding electron groups Valence Bond theory uses
hypothetical hybrid orbitals which are atomic orbitals
obtained when two or more non-equivalent orbitals of the
same atom combine in preparation for covalent bond
formation H bridisation is the term a lied to the mixin of
atomic orbitals in an atom (usually a central atom) togenerate a set of hybrid orbitals In the case of sp3 We can
generate four equivalent hybrid orbitals from the center
atom by mixing the s orbital and the three p orbitals
Explanation Energy level diagram Diagram of orbitals
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Explanation Energy level diagram Diagram of orbitals
a) Ground state
Carbon which act as thecenter atom has the valence
electron of 2s22p2
b) Excited state One of the electron from 2s
is promote to 2p orbitals -
equa energy eve
c) Hybridised state
One orbital of 2s and three
orbitals of 2p combined(hybrid) and rearrange
themselves to the shape
and orientation of a
tetrahedral shape
Molecular shape Tetrahedral
Angle between bond pair 10950
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2 sp2 hybridisation ~ When three electron groups surround
the central atom the center atom involved must prepare three
orbitals with equal energies to overlap with the three
surrounding electron groups In sp2 hybridisation three
equivalent (in terms of energy level) from the center atom by
mixing the one s orbital and the two p orbitals Using boron
trifluoride (BF3) as example sp2
hybridisation is explained
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8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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3 sp hybridisation ~ When two electron groups surround
the central atom we observe a linear shape which
means that the bonding orbitals must have a linear
orientation VB theory explains this by proposing thatmixing two nonequivalent orbitals of a central atom one
s and one p gives rise to two equivalent sp hybrid
orbitals that lie 1800 apart
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Beryllium which act as the
center atom has the valence
electron of 2s2
b) Excited state
One of the electron from 2s ispromote to 2p orbitals - equal
energy level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of a
linear shape
Molecular shape Linear
Angle between bond pair 1800
4 The concept of hybridisation is also useful to explain molecules with
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p y p
doubletriple bonds By using the concept of the direct overlapping
orbitals and side-touch lapping orbitals the formation of multiple
bonds in ethene C2H4 and ethyne C2H2 are described
a) Ethene C2H4 - Hybridisation take place for both carbon atoms in
ethene molecule is sp2 hybridisation
bull When both hybridised C is bonded together one of the hybridisedorbital overlapped directly between each other while the other two
hybridised orbitals overlapped directly with hydrogen atoms
z
atoms form a side-touch bond between each other and form anotherbond as shown in the diagram below
bull From the diagram with C=C there are two types of bond A sigma-
bond (σ-bond) is covalent bonds formed by orbitals overlapping end-
to-end with the electron density concentrated between the nuclei ofthe bonding atoms while the second type is called a pi bond (π-
bond) which is defined as a covalent bond formed by sideways
overlapping orbitals with electron density concentrated above and
below the plane of the nuclei of the bonding atoms
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital of 2s and two
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of atrigonal planar shape
Note that on unhybridised pz
orbital an electron is presence
Molecular shape Trigonal planar Angle between bond pair 1200
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b) Ethyne C H has the Lewis structure of H CequivC H which the
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b) Ethyne C2H2 has the Lewis structure of HminusCequivCminusH which the
bonding can be explain using sp hybridisation Table below
described how the hybridisation take place on each carbonatom and how the formation of triple bond occur
c) From the diagram formation of -CequivC- is due to the formation
of one sigma-bond by direct overlapping of one of the twohybrid orbitals on each C while the other two pi-bond bonds
are formed as a result of side-lapping of the each two
y z
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital from 2s and 2p
orbitals combined (hybrid) and
rearrange themselves to the
shape and orientation of a
linear shapeNote that on unhybridised py amp
pz orbitals electrons are
presence to form two
π-bonds Molecular shape Linear Angle between bond pair 1800
sp
pz py
5 Other examples of applications in valence bond theory includes the
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formation of nitrogen molecule N2 and hydrogen cyanide HCN
molecule
a) Nitrogen gas is Earths most abundant gas as it cover 78 of the
content of our air Nitrogen molecule is an inert gas thanks to its
short covalent bond and also its strong triple bond Therefore
a lot of heats are required to break the chemical bond of nitrogenbefore it can be applied in industries Using valence bond theory
the bonding of nitrogen molecule is explained in the diagram
b) So when the two hybridised nitrogen atom interacting amongeach other they hence formed a linear shape and two pi-bonds
(πminusbonds) are formed as a result of side-lapping of unhybridised
py and pz orbital respectively
N N
ExplanationEnergy level
diagramDiagram of orbitals
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diagram
a) Ground state
Both nitrogen whichact as the center atom
has the valence
electron of 2s22p3
b) Excited state One of the electron
from 2s is promote to 2p
orbitals - e ual ener
level
c) Hybridised state
One orbital of 2s and
one orbitals of 2p
combined (hybrid) and
rearrange themselves
to the shape and
orientation of a linear
shapeMolecular shape Linear
Angle between bond pair 1800
b) Whereas for hydrogen cyanide HCN both carbon atom and
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) y g y
nitrogen atom undergoes sp hybridisation in order to form a
linear structureExplanation
Energy level diagram for
carbon atom C
Energy level diagram for
nitrogen atom N
a) Ground state
Valence electronC 2s22p2
N 2s22p3
b) Excited state
One of the electron from
2s is promote to 2p
orbitals - equal energy
level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined
(hybrid) and form sp
hybridisation
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N
σσσσ
π
σσσσ
π
6 However there are a few limitation on valence bond theory such
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as when explaining the effect angle of bond-pair and bond-pair
electrons when there isare presence of lone pair electron in thecenter atom and also the difference of electronegativity
a) Bonding in ammonia NH3 and water H2O ~ Both NH3 and
H2O undergoes an arrangement similar to sp3 hybridisation
similar to that of C in methane Table below shows thehybridisation occur for nitrogen in ammonia and oxygen in water
Energy level diagram Energy level diagram
for nitrogen atom N for oxygen atom O
a) Ground state
Valence electron
N 2s22p3
O 2s2
2p4
b) Excited state
One of the electron
from 2s is promote to
2p orbitals
c) Hybridised state
One orbital of 2s and
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One orbital of 2s and
three orbitals from 2p
combined (hybrid) andrearrange themselves to
the shape and
orientation similar to that
of tetrahedral shape
Similar to arrangement in
tetrahedral
Similar to arrangement in
tetrahedral
of ammonia NH3 andwater H2O
Shape and angleShape trigonal pyramidal
Angle 1070
Shape bent
Angle 10450
Number of bond pair amp
lone pair electrons
Bond pair electron 3
Lone pair electron 1
Bond pair electron 2
Lone pair electron 2
bull The angle of bond pair - bond pair electrons in ammonia and0 0
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water are 1070 and 10450 respectively which is lesser than in
methane molecule (10950
) This can be explained by the factof the presence of lone-pair electrons in both ammonia and
water Since the lone pair - lone pair electron repulsion is
stronger than lone pair - bond pair electron repulsion than
bond pair - bond pair electron repulsion it is expected that therepulsion between lone pair - bond pair electrons in ammonia
is stronger than bond pair - bond pair electrons repulsion
hence caused the angle to squeeze to a smaller angle As
for water since there is a presence of lone pair - lone pair
electrons repulsion it results the bond pair - bond pair
electron repulsion to be much smaller hence caused the
angle to squeeze to a smaller angle Though valence bondtheory does not actually explain the hybridisation especially
when it relates to the repulsion occur involving lone pair
electron another theories shall be applied to study such
effect All these shall be discussed further in VSEPR theory
b) Effect of electronegativity and bonding angle between bond pair
- bond pair electrons ~
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- bond pair electrons
i As discussed earlier in the bonding of water (H2
O) the
molecular shape and angle is described in the diagram below
Sul hur which is also an element from Grou 16 formed h dro en
sulphide H2S when sulphur react with hydrogen However unlike
water the bonding angles is much smaller compare to water This is
due to the difference of electronegativity and also the bond length
between O and S in molecule Since O is more electronegative
compare to S the O-H bond is pulled closer toward O Furthermorethe bond length of O-H is shorter compare to S-H As a result the
bonding pair-bonding pair electrons repulse greater with each other in
H-O-H and caused the angle become greater
ii Another example is between NH3 and PH3 The orbital
f
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diagram for both NH3 and PH3 are described below Both
nitrogen N and phosphorous P are from the same groupwhich is Group 15
Similar to the above case P in phosphine is less
electronegative than N in ammonia and bond length of N-H
is shorter than P-H As a result H is pull closer to N and
repulsion between bonding of H-N-H is greater compare toH-P-H and caused the angle between H-N-H is greater
compare to H-P-H
324 Valence Shell Electron Pair Repulsion (VSEPR) Theory
1 M l l t i th th di i l t f t
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1 Molecular geometry is the three-dimensional arrangement of atoms
in a molecule A moleculersquos geometry affects its physical andchemical properties such as melting point boiling point density
and the types of reactions it undergoes The basic concept of
VSEPR theory is based on the three general rules below
a) As far as electron-pair repulsion is concerned double bonds andtriple bonds can be treated like single bonds This approximation is
good for qualitative purposes However you should realize that in
ldquo rdquo
because there are two or three bonds between two atoms theelectron density occupies more space
b) If a molecule has two or more resonance structures we can apply
the VSEPR model to any one of them Formal charges are usually
not shown
c) The order of repulsion strength of lone pair and bond pair are
lone-pair amp lone-pair electrons repulsion are the strongest
followed by lone-pair amp bond-pair electrons repulsion while
bond-pair amp bond-pair electrons repulsion is the weakest
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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CHAPTER 3 CHEMICAL BONDING
31 Ionic bonding
32 Covalent bonding
33 Metallic bonding
34 Intermolecular forces Van der Waals forces andhydrogen bonding
Topic
2007 2008 2009 2010 2011 20122013
Sem 12014
Sem 1
P1 P2 P1 P2 P1 P2 P1 P2 P1 P2 P1 P2 AB
CA
B
C
3
Chemical
Bonding
43c
6a4 5a 3
3b
c1 2
2
5b2
1b
5c
7a
319
b3 19
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INTERACTION
BETWEEN ELEMENTS
Metal and
non-metalNon - metal and
non-metal
Metal and
metal
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bull A Lewis dot symbol consists of the symbol of an element and
one dot for each valence electron in an atom of the element
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31 Ionic Bonding
bull The central idea of the ionic bonding model is the transfer of
electrons from metal atoms to non-metal atoms to form
ions that come together in a solid ionic compound where
ionic bond is formed in between oppositely charged ions
by electrostatic attraction forcesbull For example in the formation of sodium fluoride NaF
Sodium atom Fluorine atom Sodium fluoride
Electronicconfiguration
Na (1s22s22p63s1) F (1s22s22p5) Na+ F-
(1s22s22p6) (1s22s22p6)
Orbitaldiagrams Na F Na+ F-
Lewis
diagram
1s 2s 2p 3s
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bull The interaction between sodium atom and fluorine atom
occur where sodium atom (with low ionisation energy)
donates electron to fluorine (with high electron affinity) to formsodium ion Na+ and fluoride ion F- respectively Note that
both ions have achieved octet arrangement of ns2np6 as it is
the most stable form of an ion formedbull The oppositely charged Na+ and F- form a giant ionic crystal
lattice with very high melting point via electrostatic attraction
i Magnesium chloride MgCl2 ii potassium oxide K2O
iii Calcium sulphide CaS iv aluminium oxide Al2O3
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bull According to Coulumbs Law electrostatic energy between
two oppositely charged substance (A and B) is directly
proportional to the charge carried by each ions yet inversely
proportional to the distance between them
bull This relationship helps us predict trends in lattice energy and
explain the effects of ionic size and charge
( ) minus+
minus+
minus
times
prop r r
energylatticeor EnergyticElectrosta
nn
a) Effect of ionic size As we move down a group in theperiodic table the ionic radius increases Therefore the
electrostatic energy between cations and anions decreases
because the inter-ionic distance is greater thus the lattice
energies of their compounds should decrease as well This
prediction is borne out by the alkali-metal halides note the
regular decrease in lattice energy down a group whether we
hold the cation constant (LiF to LiI)
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b) Effect of ionic charge When we compare lithium fluoride
with magnesium oxide we find cations of about equal radii
(Li+ = 76 pm and Mg2+ = 72 pm) and anions of about equalradii (F- = 133 pm and O2- = 140 pm) Thus the only
significant difference is the ionic charge LiF contains the
singly charged Li+ and F- ions whereas MgO contains the
doubly charged Mg2+ and O2- ions The difference in their
lattice energies is
∆H of LiF = - 1050 kJ mol-1
∆Hlattice of MgO = - 3923 kJ mol-1
This nearly fourfold increase in ∆Hlattice reflects the fourfold
increase in the product of the charges (1 x 1 vs 2 x 2) in the
numerator of equation above
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3 Properties of Ionic Compound
a) Melting point - Ionic compound has giant ionic crystal lattice
which are hold by strong electrostatic attraction forces by
repeating of oppositely charged ions
Therefore the melting point of ionic compounds are usuallyvery high as a lot of energies are required to overcome the
electrostatic attraction forces and melted to form free moving
ions Therefore ionic compounds have very high melting
point
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b) Conductivity of electricity - Most ionic compounds do not
conduct electricity (insulator) in the solid state but do
conduct it when melted or when dissolved in water (exceptsome super-ionic conductors such as AgI which have
remarkable conductivity in the solid state) Solid ionic salt
consists of immobilized ions When it melts or dissolveshowever the ions are free to move and carry an electric
current
-
are hard (does not dent) rigid (does not bend) and brittle(cracks without deforming) These properties are due to the
powerful attractive forces that hold the ions in specific
positions throughout the crystal Moving the ions out of
position requires overcoming these forces so the sample
resists denting and bending If enough pressure is applied
ions of like charge are brought next to each other and
repulsive forces crack the ionic solid suddenly
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32 Covalent bondbull Studies of covalent bond was widely developed ever since
Lewis suggested that a chemical bond exist in a hydrogen gas
occur by sharing en electron between two hydrogen atoms
bull Electron pair that connect the 2 hydrogen atoms is called
covalent bond a bond in which two electrons are shared by
wo a oms an e e ec ron pa r a on e ween e wohydrogen atoms is also called as bonding pair electrons
bull In a covalent bond each electron in a shared pair is attracted to
the nuclei of both atoms This attraction holds the two atoms in
H2 together and is responsible for the formation of covalentbonds in other molecules
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2 A Lewis structure is a representation of covalent bonding
in which shared electron pairs are shown either as lines
or as pairs of dots between two atoms and lone pairsare shown as pairs of dots on individual atoms Only
valence electrons are shown in a Lewis structure
a) Consider the fluorine molecule F2 The electronconfiguration of F is 1s22s22p5 The 1s electrons are in
inner shell which is nearest to the nucleus For this reason
the do not artici ate in bond formation Thus each F
atom has seven valence electrons (2s22p5) Thereforeeach fluorine atom needed one electron to achieve octet
configuration (ns2np6)
F 1 s22s22p5 F 1 s22s22p5 1s2 2s2 2p6 1s2 2s2 2p6
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bull Oxygen atom has electronic configuration of 1s22s22p4 To
achieve stable octet configuration (ns2np6) each oxygen
atom need 2 electrons Hence when 2 oxygen atomsinteract they shared two electrons in between each other as
described in diagram below
O 1s2
2s2
2p4
O 1s2
2s2
2p4
1s2
2s2
2p6
1s2
2s2
2p6
bull From the structure of oxygen molecule formed each oxygen
atom shared two electrons from each other to form a double
bond in order to achieve octet configuration among each
oxygen atom
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bull Nitrogen has electronic configuration of 1s22s22p3 and
required 3 electrons to achieve octet configuration (ns2np6)
In this case each nitrogen atom shared 3 electrons fromeach of its atom to form triple bond
N 1 s22s22p3 N 1 s22s22p3 1s2 2s2 2p6 1s2 2s2 2p6
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Water H2O
H 1 s1 O 2 s22p4
Carbon dioxide CO2
C 2 s22p2 O 2 s22p4
Ammonia NH3
H 1 s1 N 2 s22p3Ethene C2H4
H 1 s1 C 2 s22p2
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Hydrogen cyanide HCN
H 1 s1 N 2 s22p3 C 2 s22p2
Ethanoic acid CH3COOH
H 1 s1 O 2 s22p4 C 2 s22p2
Tetrachloromethane CCl4C 2 s22p2 Cl 3s23p5
Ethyne C2H2
H 1 s1 C 2 s22p2
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a) Note that in ethene hydrogen cyanide and ethanoic acid all
the valence electrons are used in bonding there are no lone
pairs on the carbon atoms In fact most of the stablemolecules containing carbon do not have lone pairs on the
carbon atoms
b) Multiple bonds (double bond or triple bond) are shorter than single covalent bonds Bond length is defined as the
distance between the nuclei of two covalently bonded atoms
nitrogen triple bonds are shorter than double bonds whichin turn are shorter than single bonds The shorter multiple
bonds are also more stable than single bonds
c) Covalent bond not only exist in neutral molecule but also insome molecular ions Table below shows a few example of
molecular ions which have covalent bonds in its molecule
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Carbonate ion CO32-
C 2 s22p2 O 2 s22p4
Cyanide ion CN-
N 2 s22p3 C 2 s22p2
u p a e on 4-
S 3 s23p4 O 2 s22p4
ra e on 3-
N 2 s22p3 O 2 s22p4
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321 Exception of Octet Rules
1 From the Lewis structure sketched for sulphate ion SO42- we can
see that the center metal atom (sulphur) has more than 8electrons For this the molecule is described as molecule that can
expanded octet These molecules that have more than 8
electrons are located at Period 3 and below as these center
atoms have empty d-orbital available to expand the number ofelectrons positioned in the center atom Those center atom from
Period 2 such as C N O and F can only allocate a maximum
Example Both phosphorous and nitrogen are elements from Group15 with the valence electron of ns2np3 They can both form NCl3and PCl3 respectively when react with limited amount of chlorine
however under excess chlorine only PCl 5 can be formed but
not NCl 5 Explain the statement bolded
Solution This is due to phosphorous which is from period
3 have empty d-orbital to expand the octet However
nitrogen which is from Period 2 do not have emptyorbital and can only allocate 8 electrons in its shell
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Phosphorous pentachloride PCl5P 3 s23p3 Cl 3s23p5
Sulphur hexafluoride SF6
S 3 s23p4 F 2 s22p5
Bromine pentachloride BrCl5Cl 3s23p5 Br 4s24p5
Xenon tetrafluoride XeF4
Xe 5s25p6 F 2 s22p5
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2 There are also some stable covalent compounds which have
less than 8 electrons in their center atom (incomplete octet)
The center atoms are usually metals with great number ofvalence electrons with small atomic radius such as
beryllium boron and aluminium
BeCl2Be 2s2 Cl 3s23p5 BF3B 2 s22p1 F 2 s22p5 AlCl3Al 3s23p1 F 3 s23p5
These compound possessed stability due to their short bondlength between center atom with surround atom Furthermore
they can form resonance structure between the center atom
and surrounding atoms
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3 There are also some molecules which contain an odd
number of electrons Among the most common compounds
are nitrogen monoxide (NO) and nitrogen dioxide (NO2)
Odd-electron molecules are sometimes called radicals
Nitrogen monoxide NO
N 2 s22p3 O 2 s22p4
Nitrogen dioxide NO2
N 2 s22p3 O 2 s22p4
any ra ca s are g y reac ve e reason s t at t ere
is a tendency for the unpaired electron to form a
covalent bond with an unpaired electron on another
molecule For example when two nitrogen dioxide
molecules collide they form dinitrogen tetroxide
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322 Dative bond
1 As shown in the Lewis structure of nitrate ion one of the N-Obond is drawn as rarr The bond rarr placed is called as dative
bond (also known as coordinative bond) where dative
bond is defined as a covalent bond in which one of the atoms
donate the lone pair electrons available Although theproperties of a coordinate covalent bond do not differ from
those of a normal covalent bond because all electrons are
a e no ma er w a e r source
2 Dative bond is usually applied for these few circumstances
below
a) To assist atom molecule ion that not yet achieved octet
configuration Making use of atom which has lone pair
electrons to those which are lack of electrons
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Sulphur dioxide SO2Sulphur trioxide SO3 Carbon monoxide CO
Ozone molecule O3 Water with hydrogen ion
S
OO
O
Ammonia with hydrogen ion Ammonia with boron trifluoride
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b) Formation of dimer - In order for some compounds which
have incomplete octet to achieve stability they tend to form
dimer or polymer among themselves by using dative bondTwo of the most common examples are aluminium trichloride
and beryllium dichloride
Monomer of AlCl3Dimer of aluminium chloride Al
2
Cl6
Monomer of BeCl2 Dimer of Be2Cl4 Polymer of (BeCl2)n
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c) Formation of coordination compounds - Coordination
compounds are substances that contain at least one complex
ion a species consisting of a central metal cation (either atransition metal or a main-group metal) that is bonded to
molecules andor anions called ligands via dative
(coordinative) bond
hexaaquacopper (II) ion [Cu(H2O)6]2+ tetraamminenickel (II) ion [Ni(NH3)4]
2+
Hexacyanoferrate (III) ion [Fe(CN)6]3- Trioxalatocobaltate(III)ion [Co(C2O4)3]
3-
3 2 3 Hybridisation Theory
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323 Hybridisation Theory
Valence bond theory
bull The basic principle of valence bond theory is that a covalentbond is formed when orbitals of two atoms overlap and the
overlapped region which is between the nuclei is occupied by
a pair of electrons The central themes of valence bond theory
derive from this principle
ndash Opposing spins of the electron pair ~ Stated in Paulis
orbitals has a maximum capacity of two electrons that musthave opposite spins For example when a covalent bond is
formed in molecule of hydrogen H2 the two 1s electrons of
two H atoms occupy the overlapping 1s orbitals and have
opposite spins
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bull Maximum overlap of bonding orbitals ~ The bond strength
depends on attraction of the nuclei for the shared electrons
so the greater the orbital overlap the stronger (morestable) the bond The extent of overlap depends on the
shapes and directions of the orbitals An s orbital is spherical
but p and d orbitals have more electron density in one
direction than in another Thus whenever possible a bond
involving p or d orbitals will be oriented in the direction that
maximizes overlap For example in hydrogen fluoride (HF)
bond the 1s orbital of H overlaps the half-filled 2p orbital of Falong the long axis of that orbital Any other direction would
result in less overlap and thus a weaker bond
H b idi ti f t i bit l T t f th
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bull Hybridisation of atomic orbitals ~ To account for the
bonding in simple diatomic molecules like HF we picture the
direct overlap of s and p orbitals of isolated atoms But howcan we account for the shapes of so many molecules and
polyatomic ions through the overlap of spherical s orbitals
dumbbell-shaped p orbitals and cloverleaf-shaped d orbitals
Linus Pauling proposed that the valence atomic orbitals in the
molecule are different from those in the isolated atoms The
spatial orientations of these new orbitals lead to more stable
bonds and are consistent with observed molecular shapesThe process of orbital mixing is called hybridisation and
the new atomic orbitals are called hybrid orbitals Two key
points about the number and type of hybrid orbitals are that
i The number of hybrid orbitals obtained equals the number ofatomic orbitals mixed
ii The type of hybrid orbitals obtained varies with the types of
atomic orbitals mixed
3 2 3 1 T f h b idi ti
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3231 Type of hybridisation
1 sp3 hybridisation ~ When four electron groups surround
the central atom the center atom involved must prepare fourorbitals with equal energies to overlap with the four
surrounding electron groups Valence Bond theory uses
hypothetical hybrid orbitals which are atomic orbitals
obtained when two or more non-equivalent orbitals of the
same atom combine in preparation for covalent bond
formation H bridisation is the term a lied to the mixin of
atomic orbitals in an atom (usually a central atom) togenerate a set of hybrid orbitals In the case of sp3 We can
generate four equivalent hybrid orbitals from the center
atom by mixing the s orbital and the three p orbitals
Explanation Energy level diagram Diagram of orbitals
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Explanation Energy level diagram Diagram of orbitals
a) Ground state
Carbon which act as thecenter atom has the valence
electron of 2s22p2
b) Excited state One of the electron from 2s
is promote to 2p orbitals -
equa energy eve
c) Hybridised state
One orbital of 2s and three
orbitals of 2p combined(hybrid) and rearrange
themselves to the shape
and orientation of a
tetrahedral shape
Molecular shape Tetrahedral
Angle between bond pair 10950
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2 sp2 hybridisation ~ When three electron groups surround
the central atom the center atom involved must prepare three
orbitals with equal energies to overlap with the three
surrounding electron groups In sp2 hybridisation three
equivalent (in terms of energy level) from the center atom by
mixing the one s orbital and the two p orbitals Using boron
trifluoride (BF3) as example sp2
hybridisation is explained
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3 sp hybridisation ~ When two electron groups surround
the central atom we observe a linear shape which
means that the bonding orbitals must have a linear
orientation VB theory explains this by proposing thatmixing two nonequivalent orbitals of a central atom one
s and one p gives rise to two equivalent sp hybrid
orbitals that lie 1800 apart
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Beryllium which act as the
center atom has the valence
electron of 2s2
b) Excited state
One of the electron from 2s ispromote to 2p orbitals - equal
energy level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of a
linear shape
Molecular shape Linear
Angle between bond pair 1800
4 The concept of hybridisation is also useful to explain molecules with
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p y p
doubletriple bonds By using the concept of the direct overlapping
orbitals and side-touch lapping orbitals the formation of multiple
bonds in ethene C2H4 and ethyne C2H2 are described
a) Ethene C2H4 - Hybridisation take place for both carbon atoms in
ethene molecule is sp2 hybridisation
bull When both hybridised C is bonded together one of the hybridisedorbital overlapped directly between each other while the other two
hybridised orbitals overlapped directly with hydrogen atoms
z
atoms form a side-touch bond between each other and form anotherbond as shown in the diagram below
bull From the diagram with C=C there are two types of bond A sigma-
bond (σ-bond) is covalent bonds formed by orbitals overlapping end-
to-end with the electron density concentrated between the nuclei ofthe bonding atoms while the second type is called a pi bond (π-
bond) which is defined as a covalent bond formed by sideways
overlapping orbitals with electron density concentrated above and
below the plane of the nuclei of the bonding atoms
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital of 2s and two
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of atrigonal planar shape
Note that on unhybridised pz
orbital an electron is presence
Molecular shape Trigonal planar Angle between bond pair 1200
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b) Ethyne C H has the Lewis structure of H CequivC H which the
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b) Ethyne C2H2 has the Lewis structure of HminusCequivCminusH which the
bonding can be explain using sp hybridisation Table below
described how the hybridisation take place on each carbonatom and how the formation of triple bond occur
c) From the diagram formation of -CequivC- is due to the formation
of one sigma-bond by direct overlapping of one of the twohybrid orbitals on each C while the other two pi-bond bonds
are formed as a result of side-lapping of the each two
y z
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital from 2s and 2p
orbitals combined (hybrid) and
rearrange themselves to the
shape and orientation of a
linear shapeNote that on unhybridised py amp
pz orbitals electrons are
presence to form two
π-bonds Molecular shape Linear Angle between bond pair 1800
sp
pz py
5 Other examples of applications in valence bond theory includes the
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formation of nitrogen molecule N2 and hydrogen cyanide HCN
molecule
a) Nitrogen gas is Earths most abundant gas as it cover 78 of the
content of our air Nitrogen molecule is an inert gas thanks to its
short covalent bond and also its strong triple bond Therefore
a lot of heats are required to break the chemical bond of nitrogenbefore it can be applied in industries Using valence bond theory
the bonding of nitrogen molecule is explained in the diagram
b) So when the two hybridised nitrogen atom interacting amongeach other they hence formed a linear shape and two pi-bonds
(πminusbonds) are formed as a result of side-lapping of unhybridised
py and pz orbital respectively
N N
ExplanationEnergy level
diagramDiagram of orbitals
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diagram
a) Ground state
Both nitrogen whichact as the center atom
has the valence
electron of 2s22p3
b) Excited state One of the electron
from 2s is promote to 2p
orbitals - e ual ener
level
c) Hybridised state
One orbital of 2s and
one orbitals of 2p
combined (hybrid) and
rearrange themselves
to the shape and
orientation of a linear
shapeMolecular shape Linear
Angle between bond pair 1800
b) Whereas for hydrogen cyanide HCN both carbon atom and
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) y g y
nitrogen atom undergoes sp hybridisation in order to form a
linear structureExplanation
Energy level diagram for
carbon atom C
Energy level diagram for
nitrogen atom N
a) Ground state
Valence electronC 2s22p2
N 2s22p3
b) Excited state
One of the electron from
2s is promote to 2p
orbitals - equal energy
level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined
(hybrid) and form sp
hybridisation
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N
σσσσ
π
σσσσ
π
6 However there are a few limitation on valence bond theory such
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as when explaining the effect angle of bond-pair and bond-pair
electrons when there isare presence of lone pair electron in thecenter atom and also the difference of electronegativity
a) Bonding in ammonia NH3 and water H2O ~ Both NH3 and
H2O undergoes an arrangement similar to sp3 hybridisation
similar to that of C in methane Table below shows thehybridisation occur for nitrogen in ammonia and oxygen in water
Energy level diagram Energy level diagram
for nitrogen atom N for oxygen atom O
a) Ground state
Valence electron
N 2s22p3
O 2s2
2p4
b) Excited state
One of the electron
from 2s is promote to
2p orbitals
c) Hybridised state
One orbital of 2s and
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One orbital of 2s and
three orbitals from 2p
combined (hybrid) andrearrange themselves to
the shape and
orientation similar to that
of tetrahedral shape
Similar to arrangement in
tetrahedral
Similar to arrangement in
tetrahedral
of ammonia NH3 andwater H2O
Shape and angleShape trigonal pyramidal
Angle 1070
Shape bent
Angle 10450
Number of bond pair amp
lone pair electrons
Bond pair electron 3
Lone pair electron 1
Bond pair electron 2
Lone pair electron 2
bull The angle of bond pair - bond pair electrons in ammonia and0 0
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water are 1070 and 10450 respectively which is lesser than in
methane molecule (10950
) This can be explained by the factof the presence of lone-pair electrons in both ammonia and
water Since the lone pair - lone pair electron repulsion is
stronger than lone pair - bond pair electron repulsion than
bond pair - bond pair electron repulsion it is expected that therepulsion between lone pair - bond pair electrons in ammonia
is stronger than bond pair - bond pair electrons repulsion
hence caused the angle to squeeze to a smaller angle As
for water since there is a presence of lone pair - lone pair
electrons repulsion it results the bond pair - bond pair
electron repulsion to be much smaller hence caused the
angle to squeeze to a smaller angle Though valence bondtheory does not actually explain the hybridisation especially
when it relates to the repulsion occur involving lone pair
electron another theories shall be applied to study such
effect All these shall be discussed further in VSEPR theory
b) Effect of electronegativity and bonding angle between bond pair
- bond pair electrons ~
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- bond pair electrons
i As discussed earlier in the bonding of water (H2
O) the
molecular shape and angle is described in the diagram below
Sul hur which is also an element from Grou 16 formed h dro en
sulphide H2S when sulphur react with hydrogen However unlike
water the bonding angles is much smaller compare to water This is
due to the difference of electronegativity and also the bond length
between O and S in molecule Since O is more electronegative
compare to S the O-H bond is pulled closer toward O Furthermorethe bond length of O-H is shorter compare to S-H As a result the
bonding pair-bonding pair electrons repulse greater with each other in
H-O-H and caused the angle become greater
ii Another example is between NH3 and PH3 The orbital
f
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diagram for both NH3 and PH3 are described below Both
nitrogen N and phosphorous P are from the same groupwhich is Group 15
Similar to the above case P in phosphine is less
electronegative than N in ammonia and bond length of N-H
is shorter than P-H As a result H is pull closer to N and
repulsion between bonding of H-N-H is greater compare toH-P-H and caused the angle between H-N-H is greater
compare to H-P-H
324 Valence Shell Electron Pair Repulsion (VSEPR) Theory
1 M l l t i th th di i l t f t
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1 Molecular geometry is the three-dimensional arrangement of atoms
in a molecule A moleculersquos geometry affects its physical andchemical properties such as melting point boiling point density
and the types of reactions it undergoes The basic concept of
VSEPR theory is based on the three general rules below
a) As far as electron-pair repulsion is concerned double bonds andtriple bonds can be treated like single bonds This approximation is
good for qualitative purposes However you should realize that in
ldquo rdquo
because there are two or three bonds between two atoms theelectron density occupies more space
b) If a molecule has two or more resonance structures we can apply
the VSEPR model to any one of them Formal charges are usually
not shown
c) The order of repulsion strength of lone pair and bond pair are
lone-pair amp lone-pair electrons repulsion are the strongest
followed by lone-pair amp bond-pair electrons repulsion while
bond-pair amp bond-pair electrons repulsion is the weakest
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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INTERACTION
BETWEEN ELEMENTS
Metal and
non-metalNon - metal and
non-metal
Metal and
metal
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bull A Lewis dot symbol consists of the symbol of an element and
one dot for each valence electron in an atom of the element
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31 Ionic Bonding
bull The central idea of the ionic bonding model is the transfer of
electrons from metal atoms to non-metal atoms to form
ions that come together in a solid ionic compound where
ionic bond is formed in between oppositely charged ions
by electrostatic attraction forcesbull For example in the formation of sodium fluoride NaF
Sodium atom Fluorine atom Sodium fluoride
Electronicconfiguration
Na (1s22s22p63s1) F (1s22s22p5) Na+ F-
(1s22s22p6) (1s22s22p6)
Orbitaldiagrams Na F Na+ F-
Lewis
diagram
1s 2s 2p 3s
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bull The interaction between sodium atom and fluorine atom
occur where sodium atom (with low ionisation energy)
donates electron to fluorine (with high electron affinity) to formsodium ion Na+ and fluoride ion F- respectively Note that
both ions have achieved octet arrangement of ns2np6 as it is
the most stable form of an ion formedbull The oppositely charged Na+ and F- form a giant ionic crystal
lattice with very high melting point via electrostatic attraction
i Magnesium chloride MgCl2 ii potassium oxide K2O
iii Calcium sulphide CaS iv aluminium oxide Al2O3
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bull According to Coulumbs Law electrostatic energy between
two oppositely charged substance (A and B) is directly
proportional to the charge carried by each ions yet inversely
proportional to the distance between them
bull This relationship helps us predict trends in lattice energy and
explain the effects of ionic size and charge
( ) minus+
minus+
minus
times
prop r r
energylatticeor EnergyticElectrosta
nn
a) Effect of ionic size As we move down a group in theperiodic table the ionic radius increases Therefore the
electrostatic energy between cations and anions decreases
because the inter-ionic distance is greater thus the lattice
energies of their compounds should decrease as well This
prediction is borne out by the alkali-metal halides note the
regular decrease in lattice energy down a group whether we
hold the cation constant (LiF to LiI)
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b) Effect of ionic charge When we compare lithium fluoride
with magnesium oxide we find cations of about equal radii
(Li+ = 76 pm and Mg2+ = 72 pm) and anions of about equalradii (F- = 133 pm and O2- = 140 pm) Thus the only
significant difference is the ionic charge LiF contains the
singly charged Li+ and F- ions whereas MgO contains the
doubly charged Mg2+ and O2- ions The difference in their
lattice energies is
∆H of LiF = - 1050 kJ mol-1
∆Hlattice of MgO = - 3923 kJ mol-1
This nearly fourfold increase in ∆Hlattice reflects the fourfold
increase in the product of the charges (1 x 1 vs 2 x 2) in the
numerator of equation above
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3 Properties of Ionic Compound
a) Melting point - Ionic compound has giant ionic crystal lattice
which are hold by strong electrostatic attraction forces by
repeating of oppositely charged ions
Therefore the melting point of ionic compounds are usuallyvery high as a lot of energies are required to overcome the
electrostatic attraction forces and melted to form free moving
ions Therefore ionic compounds have very high melting
point
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b) Conductivity of electricity - Most ionic compounds do not
conduct electricity (insulator) in the solid state but do
conduct it when melted or when dissolved in water (exceptsome super-ionic conductors such as AgI which have
remarkable conductivity in the solid state) Solid ionic salt
consists of immobilized ions When it melts or dissolveshowever the ions are free to move and carry an electric
current
-
are hard (does not dent) rigid (does not bend) and brittle(cracks without deforming) These properties are due to the
powerful attractive forces that hold the ions in specific
positions throughout the crystal Moving the ions out of
position requires overcoming these forces so the sample
resists denting and bending If enough pressure is applied
ions of like charge are brought next to each other and
repulsive forces crack the ionic solid suddenly
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32 Covalent bondbull Studies of covalent bond was widely developed ever since
Lewis suggested that a chemical bond exist in a hydrogen gas
occur by sharing en electron between two hydrogen atoms
bull Electron pair that connect the 2 hydrogen atoms is called
covalent bond a bond in which two electrons are shared by
wo a oms an e e ec ron pa r a on e ween e wohydrogen atoms is also called as bonding pair electrons
bull In a covalent bond each electron in a shared pair is attracted to
the nuclei of both atoms This attraction holds the two atoms in
H2 together and is responsible for the formation of covalentbonds in other molecules
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2 A Lewis structure is a representation of covalent bonding
in which shared electron pairs are shown either as lines
or as pairs of dots between two atoms and lone pairsare shown as pairs of dots on individual atoms Only
valence electrons are shown in a Lewis structure
a) Consider the fluorine molecule F2 The electronconfiguration of F is 1s22s22p5 The 1s electrons are in
inner shell which is nearest to the nucleus For this reason
the do not artici ate in bond formation Thus each F
atom has seven valence electrons (2s22p5) Thereforeeach fluorine atom needed one electron to achieve octet
configuration (ns2np6)
F 1 s22s22p5 F 1 s22s22p5 1s2 2s2 2p6 1s2 2s2 2p6
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bull Oxygen atom has electronic configuration of 1s22s22p4 To
achieve stable octet configuration (ns2np6) each oxygen
atom need 2 electrons Hence when 2 oxygen atomsinteract they shared two electrons in between each other as
described in diagram below
O 1s2
2s2
2p4
O 1s2
2s2
2p4
1s2
2s2
2p6
1s2
2s2
2p6
bull From the structure of oxygen molecule formed each oxygen
atom shared two electrons from each other to form a double
bond in order to achieve octet configuration among each
oxygen atom
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bull Nitrogen has electronic configuration of 1s22s22p3 and
required 3 electrons to achieve octet configuration (ns2np6)
In this case each nitrogen atom shared 3 electrons fromeach of its atom to form triple bond
N 1 s22s22p3 N 1 s22s22p3 1s2 2s2 2p6 1s2 2s2 2p6
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Water H2O
H 1 s1 O 2 s22p4
Carbon dioxide CO2
C 2 s22p2 O 2 s22p4
Ammonia NH3
H 1 s1 N 2 s22p3Ethene C2H4
H 1 s1 C 2 s22p2
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Hydrogen cyanide HCN
H 1 s1 N 2 s22p3 C 2 s22p2
Ethanoic acid CH3COOH
H 1 s1 O 2 s22p4 C 2 s22p2
Tetrachloromethane CCl4C 2 s22p2 Cl 3s23p5
Ethyne C2H2
H 1 s1 C 2 s22p2
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a) Note that in ethene hydrogen cyanide and ethanoic acid all
the valence electrons are used in bonding there are no lone
pairs on the carbon atoms In fact most of the stablemolecules containing carbon do not have lone pairs on the
carbon atoms
b) Multiple bonds (double bond or triple bond) are shorter than single covalent bonds Bond length is defined as the
distance between the nuclei of two covalently bonded atoms
nitrogen triple bonds are shorter than double bonds whichin turn are shorter than single bonds The shorter multiple
bonds are also more stable than single bonds
c) Covalent bond not only exist in neutral molecule but also insome molecular ions Table below shows a few example of
molecular ions which have covalent bonds in its molecule
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Carbonate ion CO32-
C 2 s22p2 O 2 s22p4
Cyanide ion CN-
N 2 s22p3 C 2 s22p2
u p a e on 4-
S 3 s23p4 O 2 s22p4
ra e on 3-
N 2 s22p3 O 2 s22p4
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321 Exception of Octet Rules
1 From the Lewis structure sketched for sulphate ion SO42- we can
see that the center metal atom (sulphur) has more than 8electrons For this the molecule is described as molecule that can
expanded octet These molecules that have more than 8
electrons are located at Period 3 and below as these center
atoms have empty d-orbital available to expand the number ofelectrons positioned in the center atom Those center atom from
Period 2 such as C N O and F can only allocate a maximum
Example Both phosphorous and nitrogen are elements from Group15 with the valence electron of ns2np3 They can both form NCl3and PCl3 respectively when react with limited amount of chlorine
however under excess chlorine only PCl 5 can be formed but
not NCl 5 Explain the statement bolded
Solution This is due to phosphorous which is from period
3 have empty d-orbital to expand the octet However
nitrogen which is from Period 2 do not have emptyorbital and can only allocate 8 electrons in its shell
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Phosphorous pentachloride PCl5P 3 s23p3 Cl 3s23p5
Sulphur hexafluoride SF6
S 3 s23p4 F 2 s22p5
Bromine pentachloride BrCl5Cl 3s23p5 Br 4s24p5
Xenon tetrafluoride XeF4
Xe 5s25p6 F 2 s22p5
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2 There are also some stable covalent compounds which have
less than 8 electrons in their center atom (incomplete octet)
The center atoms are usually metals with great number ofvalence electrons with small atomic radius such as
beryllium boron and aluminium
BeCl2Be 2s2 Cl 3s23p5 BF3B 2 s22p1 F 2 s22p5 AlCl3Al 3s23p1 F 3 s23p5
These compound possessed stability due to their short bondlength between center atom with surround atom Furthermore
they can form resonance structure between the center atom
and surrounding atoms
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3 There are also some molecules which contain an odd
number of electrons Among the most common compounds
are nitrogen monoxide (NO) and nitrogen dioxide (NO2)
Odd-electron molecules are sometimes called radicals
Nitrogen monoxide NO
N 2 s22p3 O 2 s22p4
Nitrogen dioxide NO2
N 2 s22p3 O 2 s22p4
any ra ca s are g y reac ve e reason s t at t ere
is a tendency for the unpaired electron to form a
covalent bond with an unpaired electron on another
molecule For example when two nitrogen dioxide
molecules collide they form dinitrogen tetroxide
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322 Dative bond
1 As shown in the Lewis structure of nitrate ion one of the N-Obond is drawn as rarr The bond rarr placed is called as dative
bond (also known as coordinative bond) where dative
bond is defined as a covalent bond in which one of the atoms
donate the lone pair electrons available Although theproperties of a coordinate covalent bond do not differ from
those of a normal covalent bond because all electrons are
a e no ma er w a e r source
2 Dative bond is usually applied for these few circumstances
below
a) To assist atom molecule ion that not yet achieved octet
configuration Making use of atom which has lone pair
electrons to those which are lack of electrons
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Sulphur dioxide SO2Sulphur trioxide SO3 Carbon monoxide CO
Ozone molecule O3 Water with hydrogen ion
S
OO
O
Ammonia with hydrogen ion Ammonia with boron trifluoride
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b) Formation of dimer - In order for some compounds which
have incomplete octet to achieve stability they tend to form
dimer or polymer among themselves by using dative bondTwo of the most common examples are aluminium trichloride
and beryllium dichloride
Monomer of AlCl3Dimer of aluminium chloride Al
2
Cl6
Monomer of BeCl2 Dimer of Be2Cl4 Polymer of (BeCl2)n
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c) Formation of coordination compounds - Coordination
compounds are substances that contain at least one complex
ion a species consisting of a central metal cation (either atransition metal or a main-group metal) that is bonded to
molecules andor anions called ligands via dative
(coordinative) bond
hexaaquacopper (II) ion [Cu(H2O)6]2+ tetraamminenickel (II) ion [Ni(NH3)4]
2+
Hexacyanoferrate (III) ion [Fe(CN)6]3- Trioxalatocobaltate(III)ion [Co(C2O4)3]
3-
3 2 3 Hybridisation Theory
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323 Hybridisation Theory
Valence bond theory
bull The basic principle of valence bond theory is that a covalentbond is formed when orbitals of two atoms overlap and the
overlapped region which is between the nuclei is occupied by
a pair of electrons The central themes of valence bond theory
derive from this principle
ndash Opposing spins of the electron pair ~ Stated in Paulis
orbitals has a maximum capacity of two electrons that musthave opposite spins For example when a covalent bond is
formed in molecule of hydrogen H2 the two 1s electrons of
two H atoms occupy the overlapping 1s orbitals and have
opposite spins
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bull Maximum overlap of bonding orbitals ~ The bond strength
depends on attraction of the nuclei for the shared electrons
so the greater the orbital overlap the stronger (morestable) the bond The extent of overlap depends on the
shapes and directions of the orbitals An s orbital is spherical
but p and d orbitals have more electron density in one
direction than in another Thus whenever possible a bond
involving p or d orbitals will be oriented in the direction that
maximizes overlap For example in hydrogen fluoride (HF)
bond the 1s orbital of H overlaps the half-filled 2p orbital of Falong the long axis of that orbital Any other direction would
result in less overlap and thus a weaker bond
H b idi ti f t i bit l T t f th
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bull Hybridisation of atomic orbitals ~ To account for the
bonding in simple diatomic molecules like HF we picture the
direct overlap of s and p orbitals of isolated atoms But howcan we account for the shapes of so many molecules and
polyatomic ions through the overlap of spherical s orbitals
dumbbell-shaped p orbitals and cloverleaf-shaped d orbitals
Linus Pauling proposed that the valence atomic orbitals in the
molecule are different from those in the isolated atoms The
spatial orientations of these new orbitals lead to more stable
bonds and are consistent with observed molecular shapesThe process of orbital mixing is called hybridisation and
the new atomic orbitals are called hybrid orbitals Two key
points about the number and type of hybrid orbitals are that
i The number of hybrid orbitals obtained equals the number ofatomic orbitals mixed
ii The type of hybrid orbitals obtained varies with the types of
atomic orbitals mixed
3 2 3 1 T f h b idi ti
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3231 Type of hybridisation
1 sp3 hybridisation ~ When four electron groups surround
the central atom the center atom involved must prepare fourorbitals with equal energies to overlap with the four
surrounding electron groups Valence Bond theory uses
hypothetical hybrid orbitals which are atomic orbitals
obtained when two or more non-equivalent orbitals of the
same atom combine in preparation for covalent bond
formation H bridisation is the term a lied to the mixin of
atomic orbitals in an atom (usually a central atom) togenerate a set of hybrid orbitals In the case of sp3 We can
generate four equivalent hybrid orbitals from the center
atom by mixing the s orbital and the three p orbitals
Explanation Energy level diagram Diagram of orbitals
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Explanation Energy level diagram Diagram of orbitals
a) Ground state
Carbon which act as thecenter atom has the valence
electron of 2s22p2
b) Excited state One of the electron from 2s
is promote to 2p orbitals -
equa energy eve
c) Hybridised state
One orbital of 2s and three
orbitals of 2p combined(hybrid) and rearrange
themselves to the shape
and orientation of a
tetrahedral shape
Molecular shape Tetrahedral
Angle between bond pair 10950
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2 sp2 hybridisation ~ When three electron groups surround
the central atom the center atom involved must prepare three
orbitals with equal energies to overlap with the three
surrounding electron groups In sp2 hybridisation three
equivalent (in terms of energy level) from the center atom by
mixing the one s orbital and the two p orbitals Using boron
trifluoride (BF3) as example sp2
hybridisation is explained
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3 sp hybridisation ~ When two electron groups surround
the central atom we observe a linear shape which
means that the bonding orbitals must have a linear
orientation VB theory explains this by proposing thatmixing two nonequivalent orbitals of a central atom one
s and one p gives rise to two equivalent sp hybrid
orbitals that lie 1800 apart
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Beryllium which act as the
center atom has the valence
electron of 2s2
b) Excited state
One of the electron from 2s ispromote to 2p orbitals - equal
energy level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of a
linear shape
Molecular shape Linear
Angle between bond pair 1800
4 The concept of hybridisation is also useful to explain molecules with
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p y p
doubletriple bonds By using the concept of the direct overlapping
orbitals and side-touch lapping orbitals the formation of multiple
bonds in ethene C2H4 and ethyne C2H2 are described
a) Ethene C2H4 - Hybridisation take place for both carbon atoms in
ethene molecule is sp2 hybridisation
bull When both hybridised C is bonded together one of the hybridisedorbital overlapped directly between each other while the other two
hybridised orbitals overlapped directly with hydrogen atoms
z
atoms form a side-touch bond between each other and form anotherbond as shown in the diagram below
bull From the diagram with C=C there are two types of bond A sigma-
bond (σ-bond) is covalent bonds formed by orbitals overlapping end-
to-end with the electron density concentrated between the nuclei ofthe bonding atoms while the second type is called a pi bond (π-
bond) which is defined as a covalent bond formed by sideways
overlapping orbitals with electron density concentrated above and
below the plane of the nuclei of the bonding atoms
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital of 2s and two
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of atrigonal planar shape
Note that on unhybridised pz
orbital an electron is presence
Molecular shape Trigonal planar Angle between bond pair 1200
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b) Ethyne C H has the Lewis structure of H CequivC H which the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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b) Ethyne C2H2 has the Lewis structure of HminusCequivCminusH which the
bonding can be explain using sp hybridisation Table below
described how the hybridisation take place on each carbonatom and how the formation of triple bond occur
c) From the diagram formation of -CequivC- is due to the formation
of one sigma-bond by direct overlapping of one of the twohybrid orbitals on each C while the other two pi-bond bonds
are formed as a result of side-lapping of the each two
y z
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital from 2s and 2p
orbitals combined (hybrid) and
rearrange themselves to the
shape and orientation of a
linear shapeNote that on unhybridised py amp
pz orbitals electrons are
presence to form two
π-bonds Molecular shape Linear Angle between bond pair 1800
sp
pz py
5 Other examples of applications in valence bond theory includes the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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formation of nitrogen molecule N2 and hydrogen cyanide HCN
molecule
a) Nitrogen gas is Earths most abundant gas as it cover 78 of the
content of our air Nitrogen molecule is an inert gas thanks to its
short covalent bond and also its strong triple bond Therefore
a lot of heats are required to break the chemical bond of nitrogenbefore it can be applied in industries Using valence bond theory
the bonding of nitrogen molecule is explained in the diagram
b) So when the two hybridised nitrogen atom interacting amongeach other they hence formed a linear shape and two pi-bonds
(πminusbonds) are formed as a result of side-lapping of unhybridised
py and pz orbital respectively
N N
ExplanationEnergy level
diagramDiagram of orbitals
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diagram
a) Ground state
Both nitrogen whichact as the center atom
has the valence
electron of 2s22p3
b) Excited state One of the electron
from 2s is promote to 2p
orbitals - e ual ener
level
c) Hybridised state
One orbital of 2s and
one orbitals of 2p
combined (hybrid) and
rearrange themselves
to the shape and
orientation of a linear
shapeMolecular shape Linear
Angle between bond pair 1800
b) Whereas for hydrogen cyanide HCN both carbon atom and
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) y g y
nitrogen atom undergoes sp hybridisation in order to form a
linear structureExplanation
Energy level diagram for
carbon atom C
Energy level diagram for
nitrogen atom N
a) Ground state
Valence electronC 2s22p2
N 2s22p3
b) Excited state
One of the electron from
2s is promote to 2p
orbitals - equal energy
level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined
(hybrid) and form sp
hybridisation
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N
σσσσ
π
σσσσ
π
6 However there are a few limitation on valence bond theory such
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as when explaining the effect angle of bond-pair and bond-pair
electrons when there isare presence of lone pair electron in thecenter atom and also the difference of electronegativity
a) Bonding in ammonia NH3 and water H2O ~ Both NH3 and
H2O undergoes an arrangement similar to sp3 hybridisation
similar to that of C in methane Table below shows thehybridisation occur for nitrogen in ammonia and oxygen in water
Energy level diagram Energy level diagram
for nitrogen atom N for oxygen atom O
a) Ground state
Valence electron
N 2s22p3
O 2s2
2p4
b) Excited state
One of the electron
from 2s is promote to
2p orbitals
c) Hybridised state
One orbital of 2s and
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One orbital of 2s and
three orbitals from 2p
combined (hybrid) andrearrange themselves to
the shape and
orientation similar to that
of tetrahedral shape
Similar to arrangement in
tetrahedral
Similar to arrangement in
tetrahedral
of ammonia NH3 andwater H2O
Shape and angleShape trigonal pyramidal
Angle 1070
Shape bent
Angle 10450
Number of bond pair amp
lone pair electrons
Bond pair electron 3
Lone pair electron 1
Bond pair electron 2
Lone pair electron 2
bull The angle of bond pair - bond pair electrons in ammonia and0 0
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water are 1070 and 10450 respectively which is lesser than in
methane molecule (10950
) This can be explained by the factof the presence of lone-pair electrons in both ammonia and
water Since the lone pair - lone pair electron repulsion is
stronger than lone pair - bond pair electron repulsion than
bond pair - bond pair electron repulsion it is expected that therepulsion between lone pair - bond pair electrons in ammonia
is stronger than bond pair - bond pair electrons repulsion
hence caused the angle to squeeze to a smaller angle As
for water since there is a presence of lone pair - lone pair
electrons repulsion it results the bond pair - bond pair
electron repulsion to be much smaller hence caused the
angle to squeeze to a smaller angle Though valence bondtheory does not actually explain the hybridisation especially
when it relates to the repulsion occur involving lone pair
electron another theories shall be applied to study such
effect All these shall be discussed further in VSEPR theory
b) Effect of electronegativity and bonding angle between bond pair
- bond pair electrons ~
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- bond pair electrons
i As discussed earlier in the bonding of water (H2
O) the
molecular shape and angle is described in the diagram below
Sul hur which is also an element from Grou 16 formed h dro en
sulphide H2S when sulphur react with hydrogen However unlike
water the bonding angles is much smaller compare to water This is
due to the difference of electronegativity and also the bond length
between O and S in molecule Since O is more electronegative
compare to S the O-H bond is pulled closer toward O Furthermorethe bond length of O-H is shorter compare to S-H As a result the
bonding pair-bonding pair electrons repulse greater with each other in
H-O-H and caused the angle become greater
ii Another example is between NH3 and PH3 The orbital
f
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diagram for both NH3 and PH3 are described below Both
nitrogen N and phosphorous P are from the same groupwhich is Group 15
Similar to the above case P in phosphine is less
electronegative than N in ammonia and bond length of N-H
is shorter than P-H As a result H is pull closer to N and
repulsion between bonding of H-N-H is greater compare toH-P-H and caused the angle between H-N-H is greater
compare to H-P-H
324 Valence Shell Electron Pair Repulsion (VSEPR) Theory
1 M l l t i th th di i l t f t
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1 Molecular geometry is the three-dimensional arrangement of atoms
in a molecule A moleculersquos geometry affects its physical andchemical properties such as melting point boiling point density
and the types of reactions it undergoes The basic concept of
VSEPR theory is based on the three general rules below
a) As far as electron-pair repulsion is concerned double bonds andtriple bonds can be treated like single bonds This approximation is
good for qualitative purposes However you should realize that in
ldquo rdquo
because there are two or three bonds between two atoms theelectron density occupies more space
b) If a molecule has two or more resonance structures we can apply
the VSEPR model to any one of them Formal charges are usually
not shown
c) The order of repulsion strength of lone pair and bond pair are
lone-pair amp lone-pair electrons repulsion are the strongest
followed by lone-pair amp bond-pair electrons repulsion while
bond-pair amp bond-pair electrons repulsion is the weakest
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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bull A Lewis dot symbol consists of the symbol of an element and
one dot for each valence electron in an atom of the element
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31 Ionic Bonding
bull The central idea of the ionic bonding model is the transfer of
electrons from metal atoms to non-metal atoms to form
ions that come together in a solid ionic compound where
ionic bond is formed in between oppositely charged ions
by electrostatic attraction forcesbull For example in the formation of sodium fluoride NaF
Sodium atom Fluorine atom Sodium fluoride
Electronicconfiguration
Na (1s22s22p63s1) F (1s22s22p5) Na+ F-
(1s22s22p6) (1s22s22p6)
Orbitaldiagrams Na F Na+ F-
Lewis
diagram
1s 2s 2p 3s
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bull The interaction between sodium atom and fluorine atom
occur where sodium atom (with low ionisation energy)
donates electron to fluorine (with high electron affinity) to formsodium ion Na+ and fluoride ion F- respectively Note that
both ions have achieved octet arrangement of ns2np6 as it is
the most stable form of an ion formedbull The oppositely charged Na+ and F- form a giant ionic crystal
lattice with very high melting point via electrostatic attraction
i Magnesium chloride MgCl2 ii potassium oxide K2O
iii Calcium sulphide CaS iv aluminium oxide Al2O3
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bull According to Coulumbs Law electrostatic energy between
two oppositely charged substance (A and B) is directly
proportional to the charge carried by each ions yet inversely
proportional to the distance between them
bull This relationship helps us predict trends in lattice energy and
explain the effects of ionic size and charge
( ) minus+
minus+
minus
times
prop r r
energylatticeor EnergyticElectrosta
nn
a) Effect of ionic size As we move down a group in theperiodic table the ionic radius increases Therefore the
electrostatic energy between cations and anions decreases
because the inter-ionic distance is greater thus the lattice
energies of their compounds should decrease as well This
prediction is borne out by the alkali-metal halides note the
regular decrease in lattice energy down a group whether we
hold the cation constant (LiF to LiI)
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b) Effect of ionic charge When we compare lithium fluoride
with magnesium oxide we find cations of about equal radii
(Li+ = 76 pm and Mg2+ = 72 pm) and anions of about equalradii (F- = 133 pm and O2- = 140 pm) Thus the only
significant difference is the ionic charge LiF contains the
singly charged Li+ and F- ions whereas MgO contains the
doubly charged Mg2+ and O2- ions The difference in their
lattice energies is
∆H of LiF = - 1050 kJ mol-1
∆Hlattice of MgO = - 3923 kJ mol-1
This nearly fourfold increase in ∆Hlattice reflects the fourfold
increase in the product of the charges (1 x 1 vs 2 x 2) in the
numerator of equation above
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3 Properties of Ionic Compound
a) Melting point - Ionic compound has giant ionic crystal lattice
which are hold by strong electrostatic attraction forces by
repeating of oppositely charged ions
Therefore the melting point of ionic compounds are usuallyvery high as a lot of energies are required to overcome the
electrostatic attraction forces and melted to form free moving
ions Therefore ionic compounds have very high melting
point
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b) Conductivity of electricity - Most ionic compounds do not
conduct electricity (insulator) in the solid state but do
conduct it when melted or when dissolved in water (exceptsome super-ionic conductors such as AgI which have
remarkable conductivity in the solid state) Solid ionic salt
consists of immobilized ions When it melts or dissolveshowever the ions are free to move and carry an electric
current
-
are hard (does not dent) rigid (does not bend) and brittle(cracks without deforming) These properties are due to the
powerful attractive forces that hold the ions in specific
positions throughout the crystal Moving the ions out of
position requires overcoming these forces so the sample
resists denting and bending If enough pressure is applied
ions of like charge are brought next to each other and
repulsive forces crack the ionic solid suddenly
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32 Covalent bondbull Studies of covalent bond was widely developed ever since
Lewis suggested that a chemical bond exist in a hydrogen gas
occur by sharing en electron between two hydrogen atoms
bull Electron pair that connect the 2 hydrogen atoms is called
covalent bond a bond in which two electrons are shared by
wo a oms an e e ec ron pa r a on e ween e wohydrogen atoms is also called as bonding pair electrons
bull In a covalent bond each electron in a shared pair is attracted to
the nuclei of both atoms This attraction holds the two atoms in
H2 together and is responsible for the formation of covalentbonds in other molecules
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2 A Lewis structure is a representation of covalent bonding
in which shared electron pairs are shown either as lines
or as pairs of dots between two atoms and lone pairsare shown as pairs of dots on individual atoms Only
valence electrons are shown in a Lewis structure
a) Consider the fluorine molecule F2 The electronconfiguration of F is 1s22s22p5 The 1s electrons are in
inner shell which is nearest to the nucleus For this reason
the do not artici ate in bond formation Thus each F
atom has seven valence electrons (2s22p5) Thereforeeach fluorine atom needed one electron to achieve octet
configuration (ns2np6)
F 1 s22s22p5 F 1 s22s22p5 1s2 2s2 2p6 1s2 2s2 2p6
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bull Oxygen atom has electronic configuration of 1s22s22p4 To
achieve stable octet configuration (ns2np6) each oxygen
atom need 2 electrons Hence when 2 oxygen atomsinteract they shared two electrons in between each other as
described in diagram below
O 1s2
2s2
2p4
O 1s2
2s2
2p4
1s2
2s2
2p6
1s2
2s2
2p6
bull From the structure of oxygen molecule formed each oxygen
atom shared two electrons from each other to form a double
bond in order to achieve octet configuration among each
oxygen atom
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bull Nitrogen has electronic configuration of 1s22s22p3 and
required 3 electrons to achieve octet configuration (ns2np6)
In this case each nitrogen atom shared 3 electrons fromeach of its atom to form triple bond
N 1 s22s22p3 N 1 s22s22p3 1s2 2s2 2p6 1s2 2s2 2p6
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Water H2O
H 1 s1 O 2 s22p4
Carbon dioxide CO2
C 2 s22p2 O 2 s22p4
Ammonia NH3
H 1 s1 N 2 s22p3Ethene C2H4
H 1 s1 C 2 s22p2
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Hydrogen cyanide HCN
H 1 s1 N 2 s22p3 C 2 s22p2
Ethanoic acid CH3COOH
H 1 s1 O 2 s22p4 C 2 s22p2
Tetrachloromethane CCl4C 2 s22p2 Cl 3s23p5
Ethyne C2H2
H 1 s1 C 2 s22p2
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a) Note that in ethene hydrogen cyanide and ethanoic acid all
the valence electrons are used in bonding there are no lone
pairs on the carbon atoms In fact most of the stablemolecules containing carbon do not have lone pairs on the
carbon atoms
b) Multiple bonds (double bond or triple bond) are shorter than single covalent bonds Bond length is defined as the
distance between the nuclei of two covalently bonded atoms
nitrogen triple bonds are shorter than double bonds whichin turn are shorter than single bonds The shorter multiple
bonds are also more stable than single bonds
c) Covalent bond not only exist in neutral molecule but also insome molecular ions Table below shows a few example of
molecular ions which have covalent bonds in its molecule
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Carbonate ion CO32-
C 2 s22p2 O 2 s22p4
Cyanide ion CN-
N 2 s22p3 C 2 s22p2
u p a e on 4-
S 3 s23p4 O 2 s22p4
ra e on 3-
N 2 s22p3 O 2 s22p4
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321 Exception of Octet Rules
1 From the Lewis structure sketched for sulphate ion SO42- we can
see that the center metal atom (sulphur) has more than 8electrons For this the molecule is described as molecule that can
expanded octet These molecules that have more than 8
electrons are located at Period 3 and below as these center
atoms have empty d-orbital available to expand the number ofelectrons positioned in the center atom Those center atom from
Period 2 such as C N O and F can only allocate a maximum
Example Both phosphorous and nitrogen are elements from Group15 with the valence electron of ns2np3 They can both form NCl3and PCl3 respectively when react with limited amount of chlorine
however under excess chlorine only PCl 5 can be formed but
not NCl 5 Explain the statement bolded
Solution This is due to phosphorous which is from period
3 have empty d-orbital to expand the octet However
nitrogen which is from Period 2 do not have emptyorbital and can only allocate 8 electrons in its shell
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Phosphorous pentachloride PCl5P 3 s23p3 Cl 3s23p5
Sulphur hexafluoride SF6
S 3 s23p4 F 2 s22p5
Bromine pentachloride BrCl5Cl 3s23p5 Br 4s24p5
Xenon tetrafluoride XeF4
Xe 5s25p6 F 2 s22p5
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2 There are also some stable covalent compounds which have
less than 8 electrons in their center atom (incomplete octet)
The center atoms are usually metals with great number ofvalence electrons with small atomic radius such as
beryllium boron and aluminium
BeCl2Be 2s2 Cl 3s23p5 BF3B 2 s22p1 F 2 s22p5 AlCl3Al 3s23p1 F 3 s23p5
These compound possessed stability due to their short bondlength between center atom with surround atom Furthermore
they can form resonance structure between the center atom
and surrounding atoms
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3 There are also some molecules which contain an odd
number of electrons Among the most common compounds
are nitrogen monoxide (NO) and nitrogen dioxide (NO2)
Odd-electron molecules are sometimes called radicals
Nitrogen monoxide NO
N 2 s22p3 O 2 s22p4
Nitrogen dioxide NO2
N 2 s22p3 O 2 s22p4
any ra ca s are g y reac ve e reason s t at t ere
is a tendency for the unpaired electron to form a
covalent bond with an unpaired electron on another
molecule For example when two nitrogen dioxide
molecules collide they form dinitrogen tetroxide
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322 Dative bond
1 As shown in the Lewis structure of nitrate ion one of the N-Obond is drawn as rarr The bond rarr placed is called as dative
bond (also known as coordinative bond) where dative
bond is defined as a covalent bond in which one of the atoms
donate the lone pair electrons available Although theproperties of a coordinate covalent bond do not differ from
those of a normal covalent bond because all electrons are
a e no ma er w a e r source
2 Dative bond is usually applied for these few circumstances
below
a) To assist atom molecule ion that not yet achieved octet
configuration Making use of atom which has lone pair
electrons to those which are lack of electrons
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Sulphur dioxide SO2Sulphur trioxide SO3 Carbon monoxide CO
Ozone molecule O3 Water with hydrogen ion
S
OO
O
Ammonia with hydrogen ion Ammonia with boron trifluoride
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b) Formation of dimer - In order for some compounds which
have incomplete octet to achieve stability they tend to form
dimer or polymer among themselves by using dative bondTwo of the most common examples are aluminium trichloride
and beryllium dichloride
Monomer of AlCl3Dimer of aluminium chloride Al
2
Cl6
Monomer of BeCl2 Dimer of Be2Cl4 Polymer of (BeCl2)n
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c) Formation of coordination compounds - Coordination
compounds are substances that contain at least one complex
ion a species consisting of a central metal cation (either atransition metal or a main-group metal) that is bonded to
molecules andor anions called ligands via dative
(coordinative) bond
hexaaquacopper (II) ion [Cu(H2O)6]2+ tetraamminenickel (II) ion [Ni(NH3)4]
2+
Hexacyanoferrate (III) ion [Fe(CN)6]3- Trioxalatocobaltate(III)ion [Co(C2O4)3]
3-
3 2 3 Hybridisation Theory
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323 Hybridisation Theory
Valence bond theory
bull The basic principle of valence bond theory is that a covalentbond is formed when orbitals of two atoms overlap and the
overlapped region which is between the nuclei is occupied by
a pair of electrons The central themes of valence bond theory
derive from this principle
ndash Opposing spins of the electron pair ~ Stated in Paulis
orbitals has a maximum capacity of two electrons that musthave opposite spins For example when a covalent bond is
formed in molecule of hydrogen H2 the two 1s electrons of
two H atoms occupy the overlapping 1s orbitals and have
opposite spins
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bull Maximum overlap of bonding orbitals ~ The bond strength
depends on attraction of the nuclei for the shared electrons
so the greater the orbital overlap the stronger (morestable) the bond The extent of overlap depends on the
shapes and directions of the orbitals An s orbital is spherical
but p and d orbitals have more electron density in one
direction than in another Thus whenever possible a bond
involving p or d orbitals will be oriented in the direction that
maximizes overlap For example in hydrogen fluoride (HF)
bond the 1s orbital of H overlaps the half-filled 2p orbital of Falong the long axis of that orbital Any other direction would
result in less overlap and thus a weaker bond
H b idi ti f t i bit l T t f th
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bull Hybridisation of atomic orbitals ~ To account for the
bonding in simple diatomic molecules like HF we picture the
direct overlap of s and p orbitals of isolated atoms But howcan we account for the shapes of so many molecules and
polyatomic ions through the overlap of spherical s orbitals
dumbbell-shaped p orbitals and cloverleaf-shaped d orbitals
Linus Pauling proposed that the valence atomic orbitals in the
molecule are different from those in the isolated atoms The
spatial orientations of these new orbitals lead to more stable
bonds and are consistent with observed molecular shapesThe process of orbital mixing is called hybridisation and
the new atomic orbitals are called hybrid orbitals Two key
points about the number and type of hybrid orbitals are that
i The number of hybrid orbitals obtained equals the number ofatomic orbitals mixed
ii The type of hybrid orbitals obtained varies with the types of
atomic orbitals mixed
3 2 3 1 T f h b idi ti
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3231 Type of hybridisation
1 sp3 hybridisation ~ When four electron groups surround
the central atom the center atom involved must prepare fourorbitals with equal energies to overlap with the four
surrounding electron groups Valence Bond theory uses
hypothetical hybrid orbitals which are atomic orbitals
obtained when two or more non-equivalent orbitals of the
same atom combine in preparation for covalent bond
formation H bridisation is the term a lied to the mixin of
atomic orbitals in an atom (usually a central atom) togenerate a set of hybrid orbitals In the case of sp3 We can
generate four equivalent hybrid orbitals from the center
atom by mixing the s orbital and the three p orbitals
Explanation Energy level diagram Diagram of orbitals
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Explanation Energy level diagram Diagram of orbitals
a) Ground state
Carbon which act as thecenter atom has the valence
electron of 2s22p2
b) Excited state One of the electron from 2s
is promote to 2p orbitals -
equa energy eve
c) Hybridised state
One orbital of 2s and three
orbitals of 2p combined(hybrid) and rearrange
themselves to the shape
and orientation of a
tetrahedral shape
Molecular shape Tetrahedral
Angle between bond pair 10950
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2 sp2 hybridisation ~ When three electron groups surround
the central atom the center atom involved must prepare three
orbitals with equal energies to overlap with the three
surrounding electron groups In sp2 hybridisation three
equivalent (in terms of energy level) from the center atom by
mixing the one s orbital and the two p orbitals Using boron
trifluoride (BF3) as example sp2
hybridisation is explained
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3 sp hybridisation ~ When two electron groups surround
the central atom we observe a linear shape which
means that the bonding orbitals must have a linear
orientation VB theory explains this by proposing thatmixing two nonequivalent orbitals of a central atom one
s and one p gives rise to two equivalent sp hybrid
orbitals that lie 1800 apart
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Beryllium which act as the
center atom has the valence
electron of 2s2
b) Excited state
One of the electron from 2s ispromote to 2p orbitals - equal
energy level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of a
linear shape
Molecular shape Linear
Angle between bond pair 1800
4 The concept of hybridisation is also useful to explain molecules with
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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p y p
doubletriple bonds By using the concept of the direct overlapping
orbitals and side-touch lapping orbitals the formation of multiple
bonds in ethene C2H4 and ethyne C2H2 are described
a) Ethene C2H4 - Hybridisation take place for both carbon atoms in
ethene molecule is sp2 hybridisation
bull When both hybridised C is bonded together one of the hybridisedorbital overlapped directly between each other while the other two
hybridised orbitals overlapped directly with hydrogen atoms
z
atoms form a side-touch bond between each other and form anotherbond as shown in the diagram below
bull From the diagram with C=C there are two types of bond A sigma-
bond (σ-bond) is covalent bonds formed by orbitals overlapping end-
to-end with the electron density concentrated between the nuclei ofthe bonding atoms while the second type is called a pi bond (π-
bond) which is defined as a covalent bond formed by sideways
overlapping orbitals with electron density concentrated above and
below the plane of the nuclei of the bonding atoms
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital of 2s and two
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of atrigonal planar shape
Note that on unhybridised pz
orbital an electron is presence
Molecular shape Trigonal planar Angle between bond pair 1200
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b) Ethyne C H has the Lewis structure of H CequivC H which the
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b) Ethyne C2H2 has the Lewis structure of HminusCequivCminusH which the
bonding can be explain using sp hybridisation Table below
described how the hybridisation take place on each carbonatom and how the formation of triple bond occur
c) From the diagram formation of -CequivC- is due to the formation
of one sigma-bond by direct overlapping of one of the twohybrid orbitals on each C while the other two pi-bond bonds
are formed as a result of side-lapping of the each two
y z
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital from 2s and 2p
orbitals combined (hybrid) and
rearrange themselves to the
shape and orientation of a
linear shapeNote that on unhybridised py amp
pz orbitals electrons are
presence to form two
π-bonds Molecular shape Linear Angle between bond pair 1800
sp
pz py
5 Other examples of applications in valence bond theory includes the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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formation of nitrogen molecule N2 and hydrogen cyanide HCN
molecule
a) Nitrogen gas is Earths most abundant gas as it cover 78 of the
content of our air Nitrogen molecule is an inert gas thanks to its
short covalent bond and also its strong triple bond Therefore
a lot of heats are required to break the chemical bond of nitrogenbefore it can be applied in industries Using valence bond theory
the bonding of nitrogen molecule is explained in the diagram
b) So when the two hybridised nitrogen atom interacting amongeach other they hence formed a linear shape and two pi-bonds
(πminusbonds) are formed as a result of side-lapping of unhybridised
py and pz orbital respectively
N N
ExplanationEnergy level
diagramDiagram of orbitals
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diagram
a) Ground state
Both nitrogen whichact as the center atom
has the valence
electron of 2s22p3
b) Excited state One of the electron
from 2s is promote to 2p
orbitals - e ual ener
level
c) Hybridised state
One orbital of 2s and
one orbitals of 2p
combined (hybrid) and
rearrange themselves
to the shape and
orientation of a linear
shapeMolecular shape Linear
Angle between bond pair 1800
b) Whereas for hydrogen cyanide HCN both carbon atom and
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) y g y
nitrogen atom undergoes sp hybridisation in order to form a
linear structureExplanation
Energy level diagram for
carbon atom C
Energy level diagram for
nitrogen atom N
a) Ground state
Valence electronC 2s22p2
N 2s22p3
b) Excited state
One of the electron from
2s is promote to 2p
orbitals - equal energy
level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined
(hybrid) and form sp
hybridisation
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N
σσσσ
π
σσσσ
π
6 However there are a few limitation on valence bond theory such
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as when explaining the effect angle of bond-pair and bond-pair
electrons when there isare presence of lone pair electron in thecenter atom and also the difference of electronegativity
a) Bonding in ammonia NH3 and water H2O ~ Both NH3 and
H2O undergoes an arrangement similar to sp3 hybridisation
similar to that of C in methane Table below shows thehybridisation occur for nitrogen in ammonia and oxygen in water
Energy level diagram Energy level diagram
for nitrogen atom N for oxygen atom O
a) Ground state
Valence electron
N 2s22p3
O 2s2
2p4
b) Excited state
One of the electron
from 2s is promote to
2p orbitals
c) Hybridised state
One orbital of 2s and
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One orbital of 2s and
three orbitals from 2p
combined (hybrid) andrearrange themselves to
the shape and
orientation similar to that
of tetrahedral shape
Similar to arrangement in
tetrahedral
Similar to arrangement in
tetrahedral
of ammonia NH3 andwater H2O
Shape and angleShape trigonal pyramidal
Angle 1070
Shape bent
Angle 10450
Number of bond pair amp
lone pair electrons
Bond pair electron 3
Lone pair electron 1
Bond pair electron 2
Lone pair electron 2
bull The angle of bond pair - bond pair electrons in ammonia and0 0
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water are 1070 and 10450 respectively which is lesser than in
methane molecule (10950
) This can be explained by the factof the presence of lone-pair electrons in both ammonia and
water Since the lone pair - lone pair electron repulsion is
stronger than lone pair - bond pair electron repulsion than
bond pair - bond pair electron repulsion it is expected that therepulsion between lone pair - bond pair electrons in ammonia
is stronger than bond pair - bond pair electrons repulsion
hence caused the angle to squeeze to a smaller angle As
for water since there is a presence of lone pair - lone pair
electrons repulsion it results the bond pair - bond pair
electron repulsion to be much smaller hence caused the
angle to squeeze to a smaller angle Though valence bondtheory does not actually explain the hybridisation especially
when it relates to the repulsion occur involving lone pair
electron another theories shall be applied to study such
effect All these shall be discussed further in VSEPR theory
b) Effect of electronegativity and bonding angle between bond pair
- bond pair electrons ~
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- bond pair electrons
i As discussed earlier in the bonding of water (H2
O) the
molecular shape and angle is described in the diagram below
Sul hur which is also an element from Grou 16 formed h dro en
sulphide H2S when sulphur react with hydrogen However unlike
water the bonding angles is much smaller compare to water This is
due to the difference of electronegativity and also the bond length
between O and S in molecule Since O is more electronegative
compare to S the O-H bond is pulled closer toward O Furthermorethe bond length of O-H is shorter compare to S-H As a result the
bonding pair-bonding pair electrons repulse greater with each other in
H-O-H and caused the angle become greater
ii Another example is between NH3 and PH3 The orbital
f
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diagram for both NH3 and PH3 are described below Both
nitrogen N and phosphorous P are from the same groupwhich is Group 15
Similar to the above case P in phosphine is less
electronegative than N in ammonia and bond length of N-H
is shorter than P-H As a result H is pull closer to N and
repulsion between bonding of H-N-H is greater compare toH-P-H and caused the angle between H-N-H is greater
compare to H-P-H
324 Valence Shell Electron Pair Repulsion (VSEPR) Theory
1 M l l t i th th di i l t f t
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1 Molecular geometry is the three-dimensional arrangement of atoms
in a molecule A moleculersquos geometry affects its physical andchemical properties such as melting point boiling point density
and the types of reactions it undergoes The basic concept of
VSEPR theory is based on the three general rules below
a) As far as electron-pair repulsion is concerned double bonds andtriple bonds can be treated like single bonds This approximation is
good for qualitative purposes However you should realize that in
ldquo rdquo
because there are two or three bonds between two atoms theelectron density occupies more space
b) If a molecule has two or more resonance structures we can apply
the VSEPR model to any one of them Formal charges are usually
not shown
c) The order of repulsion strength of lone pair and bond pair are
lone-pair amp lone-pair electrons repulsion are the strongest
followed by lone-pair amp bond-pair electrons repulsion while
bond-pair amp bond-pair electrons repulsion is the weakest
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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31 Ionic Bonding
bull The central idea of the ionic bonding model is the transfer of
electrons from metal atoms to non-metal atoms to form
ions that come together in a solid ionic compound where
ionic bond is formed in between oppositely charged ions
by electrostatic attraction forcesbull For example in the formation of sodium fluoride NaF
Sodium atom Fluorine atom Sodium fluoride
Electronicconfiguration
Na (1s22s22p63s1) F (1s22s22p5) Na+ F-
(1s22s22p6) (1s22s22p6)
Orbitaldiagrams Na F Na+ F-
Lewis
diagram
1s 2s 2p 3s
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bull The interaction between sodium atom and fluorine atom
occur where sodium atom (with low ionisation energy)
donates electron to fluorine (with high electron affinity) to formsodium ion Na+ and fluoride ion F- respectively Note that
both ions have achieved octet arrangement of ns2np6 as it is
the most stable form of an ion formedbull The oppositely charged Na+ and F- form a giant ionic crystal
lattice with very high melting point via electrostatic attraction
i Magnesium chloride MgCl2 ii potassium oxide K2O
iii Calcium sulphide CaS iv aluminium oxide Al2O3
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bull According to Coulumbs Law electrostatic energy between
two oppositely charged substance (A and B) is directly
proportional to the charge carried by each ions yet inversely
proportional to the distance between them
bull This relationship helps us predict trends in lattice energy and
explain the effects of ionic size and charge
( ) minus+
minus+
minus
times
prop r r
energylatticeor EnergyticElectrosta
nn
a) Effect of ionic size As we move down a group in theperiodic table the ionic radius increases Therefore the
electrostatic energy between cations and anions decreases
because the inter-ionic distance is greater thus the lattice
energies of their compounds should decrease as well This
prediction is borne out by the alkali-metal halides note the
regular decrease in lattice energy down a group whether we
hold the cation constant (LiF to LiI)
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b) Effect of ionic charge When we compare lithium fluoride
with magnesium oxide we find cations of about equal radii
(Li+ = 76 pm and Mg2+ = 72 pm) and anions of about equalradii (F- = 133 pm and O2- = 140 pm) Thus the only
significant difference is the ionic charge LiF contains the
singly charged Li+ and F- ions whereas MgO contains the
doubly charged Mg2+ and O2- ions The difference in their
lattice energies is
∆H of LiF = - 1050 kJ mol-1
∆Hlattice of MgO = - 3923 kJ mol-1
This nearly fourfold increase in ∆Hlattice reflects the fourfold
increase in the product of the charges (1 x 1 vs 2 x 2) in the
numerator of equation above
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3 Properties of Ionic Compound
a) Melting point - Ionic compound has giant ionic crystal lattice
which are hold by strong electrostatic attraction forces by
repeating of oppositely charged ions
Therefore the melting point of ionic compounds are usuallyvery high as a lot of energies are required to overcome the
electrostatic attraction forces and melted to form free moving
ions Therefore ionic compounds have very high melting
point
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b) Conductivity of electricity - Most ionic compounds do not
conduct electricity (insulator) in the solid state but do
conduct it when melted or when dissolved in water (exceptsome super-ionic conductors such as AgI which have
remarkable conductivity in the solid state) Solid ionic salt
consists of immobilized ions When it melts or dissolveshowever the ions are free to move and carry an electric
current
-
are hard (does not dent) rigid (does not bend) and brittle(cracks without deforming) These properties are due to the
powerful attractive forces that hold the ions in specific
positions throughout the crystal Moving the ions out of
position requires overcoming these forces so the sample
resists denting and bending If enough pressure is applied
ions of like charge are brought next to each other and
repulsive forces crack the ionic solid suddenly
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32 Covalent bondbull Studies of covalent bond was widely developed ever since
Lewis suggested that a chemical bond exist in a hydrogen gas
occur by sharing en electron between two hydrogen atoms
bull Electron pair that connect the 2 hydrogen atoms is called
covalent bond a bond in which two electrons are shared by
wo a oms an e e ec ron pa r a on e ween e wohydrogen atoms is also called as bonding pair electrons
bull In a covalent bond each electron in a shared pair is attracted to
the nuclei of both atoms This attraction holds the two atoms in
H2 together and is responsible for the formation of covalentbonds in other molecules
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2 A Lewis structure is a representation of covalent bonding
in which shared electron pairs are shown either as lines
or as pairs of dots between two atoms and lone pairsare shown as pairs of dots on individual atoms Only
valence electrons are shown in a Lewis structure
a) Consider the fluorine molecule F2 The electronconfiguration of F is 1s22s22p5 The 1s electrons are in
inner shell which is nearest to the nucleus For this reason
the do not artici ate in bond formation Thus each F
atom has seven valence electrons (2s22p5) Thereforeeach fluorine atom needed one electron to achieve octet
configuration (ns2np6)
F 1 s22s22p5 F 1 s22s22p5 1s2 2s2 2p6 1s2 2s2 2p6
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bull Oxygen atom has electronic configuration of 1s22s22p4 To
achieve stable octet configuration (ns2np6) each oxygen
atom need 2 electrons Hence when 2 oxygen atomsinteract they shared two electrons in between each other as
described in diagram below
O 1s2
2s2
2p4
O 1s2
2s2
2p4
1s2
2s2
2p6
1s2
2s2
2p6
bull From the structure of oxygen molecule formed each oxygen
atom shared two electrons from each other to form a double
bond in order to achieve octet configuration among each
oxygen atom
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bull Nitrogen has electronic configuration of 1s22s22p3 and
required 3 electrons to achieve octet configuration (ns2np6)
In this case each nitrogen atom shared 3 electrons fromeach of its atom to form triple bond
N 1 s22s22p3 N 1 s22s22p3 1s2 2s2 2p6 1s2 2s2 2p6
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Water H2O
H 1 s1 O 2 s22p4
Carbon dioxide CO2
C 2 s22p2 O 2 s22p4
Ammonia NH3
H 1 s1 N 2 s22p3Ethene C2H4
H 1 s1 C 2 s22p2
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Hydrogen cyanide HCN
H 1 s1 N 2 s22p3 C 2 s22p2
Ethanoic acid CH3COOH
H 1 s1 O 2 s22p4 C 2 s22p2
Tetrachloromethane CCl4C 2 s22p2 Cl 3s23p5
Ethyne C2H2
H 1 s1 C 2 s22p2
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a) Note that in ethene hydrogen cyanide and ethanoic acid all
the valence electrons are used in bonding there are no lone
pairs on the carbon atoms In fact most of the stablemolecules containing carbon do not have lone pairs on the
carbon atoms
b) Multiple bonds (double bond or triple bond) are shorter than single covalent bonds Bond length is defined as the
distance between the nuclei of two covalently bonded atoms
nitrogen triple bonds are shorter than double bonds whichin turn are shorter than single bonds The shorter multiple
bonds are also more stable than single bonds
c) Covalent bond not only exist in neutral molecule but also insome molecular ions Table below shows a few example of
molecular ions which have covalent bonds in its molecule
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Carbonate ion CO32-
C 2 s22p2 O 2 s22p4
Cyanide ion CN-
N 2 s22p3 C 2 s22p2
u p a e on 4-
S 3 s23p4 O 2 s22p4
ra e on 3-
N 2 s22p3 O 2 s22p4
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321 Exception of Octet Rules
1 From the Lewis structure sketched for sulphate ion SO42- we can
see that the center metal atom (sulphur) has more than 8electrons For this the molecule is described as molecule that can
expanded octet These molecules that have more than 8
electrons are located at Period 3 and below as these center
atoms have empty d-orbital available to expand the number ofelectrons positioned in the center atom Those center atom from
Period 2 such as C N O and F can only allocate a maximum
Example Both phosphorous and nitrogen are elements from Group15 with the valence electron of ns2np3 They can both form NCl3and PCl3 respectively when react with limited amount of chlorine
however under excess chlorine only PCl 5 can be formed but
not NCl 5 Explain the statement bolded
Solution This is due to phosphorous which is from period
3 have empty d-orbital to expand the octet However
nitrogen which is from Period 2 do not have emptyorbital and can only allocate 8 electrons in its shell
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Phosphorous pentachloride PCl5P 3 s23p3 Cl 3s23p5
Sulphur hexafluoride SF6
S 3 s23p4 F 2 s22p5
Bromine pentachloride BrCl5Cl 3s23p5 Br 4s24p5
Xenon tetrafluoride XeF4
Xe 5s25p6 F 2 s22p5
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2 There are also some stable covalent compounds which have
less than 8 electrons in their center atom (incomplete octet)
The center atoms are usually metals with great number ofvalence electrons with small atomic radius such as
beryllium boron and aluminium
BeCl2Be 2s2 Cl 3s23p5 BF3B 2 s22p1 F 2 s22p5 AlCl3Al 3s23p1 F 3 s23p5
These compound possessed stability due to their short bondlength between center atom with surround atom Furthermore
they can form resonance structure between the center atom
and surrounding atoms
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3 There are also some molecules which contain an odd
number of electrons Among the most common compounds
are nitrogen monoxide (NO) and nitrogen dioxide (NO2)
Odd-electron molecules are sometimes called radicals
Nitrogen monoxide NO
N 2 s22p3 O 2 s22p4
Nitrogen dioxide NO2
N 2 s22p3 O 2 s22p4
any ra ca s are g y reac ve e reason s t at t ere
is a tendency for the unpaired electron to form a
covalent bond with an unpaired electron on another
molecule For example when two nitrogen dioxide
molecules collide they form dinitrogen tetroxide
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322 Dative bond
1 As shown in the Lewis structure of nitrate ion one of the N-Obond is drawn as rarr The bond rarr placed is called as dative
bond (also known as coordinative bond) where dative
bond is defined as a covalent bond in which one of the atoms
donate the lone pair electrons available Although theproperties of a coordinate covalent bond do not differ from
those of a normal covalent bond because all electrons are
a e no ma er w a e r source
2 Dative bond is usually applied for these few circumstances
below
a) To assist atom molecule ion that not yet achieved octet
configuration Making use of atom which has lone pair
electrons to those which are lack of electrons
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Sulphur dioxide SO2Sulphur trioxide SO3 Carbon monoxide CO
Ozone molecule O3 Water with hydrogen ion
S
OO
O
Ammonia with hydrogen ion Ammonia with boron trifluoride
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b) Formation of dimer - In order for some compounds which
have incomplete octet to achieve stability they tend to form
dimer or polymer among themselves by using dative bondTwo of the most common examples are aluminium trichloride
and beryllium dichloride
Monomer of AlCl3Dimer of aluminium chloride Al
2
Cl6
Monomer of BeCl2 Dimer of Be2Cl4 Polymer of (BeCl2)n
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c) Formation of coordination compounds - Coordination
compounds are substances that contain at least one complex
ion a species consisting of a central metal cation (either atransition metal or a main-group metal) that is bonded to
molecules andor anions called ligands via dative
(coordinative) bond
hexaaquacopper (II) ion [Cu(H2O)6]2+ tetraamminenickel (II) ion [Ni(NH3)4]
2+
Hexacyanoferrate (III) ion [Fe(CN)6]3- Trioxalatocobaltate(III)ion [Co(C2O4)3]
3-
3 2 3 Hybridisation Theory
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323 Hybridisation Theory
Valence bond theory
bull The basic principle of valence bond theory is that a covalentbond is formed when orbitals of two atoms overlap and the
overlapped region which is between the nuclei is occupied by
a pair of electrons The central themes of valence bond theory
derive from this principle
ndash Opposing spins of the electron pair ~ Stated in Paulis
orbitals has a maximum capacity of two electrons that musthave opposite spins For example when a covalent bond is
formed in molecule of hydrogen H2 the two 1s electrons of
two H atoms occupy the overlapping 1s orbitals and have
opposite spins
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bull Maximum overlap of bonding orbitals ~ The bond strength
depends on attraction of the nuclei for the shared electrons
so the greater the orbital overlap the stronger (morestable) the bond The extent of overlap depends on the
shapes and directions of the orbitals An s orbital is spherical
but p and d orbitals have more electron density in one
direction than in another Thus whenever possible a bond
involving p or d orbitals will be oriented in the direction that
maximizes overlap For example in hydrogen fluoride (HF)
bond the 1s orbital of H overlaps the half-filled 2p orbital of Falong the long axis of that orbital Any other direction would
result in less overlap and thus a weaker bond
H b idi ti f t i bit l T t f th
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bull Hybridisation of atomic orbitals ~ To account for the
bonding in simple diatomic molecules like HF we picture the
direct overlap of s and p orbitals of isolated atoms But howcan we account for the shapes of so many molecules and
polyatomic ions through the overlap of spherical s orbitals
dumbbell-shaped p orbitals and cloverleaf-shaped d orbitals
Linus Pauling proposed that the valence atomic orbitals in the
molecule are different from those in the isolated atoms The
spatial orientations of these new orbitals lead to more stable
bonds and are consistent with observed molecular shapesThe process of orbital mixing is called hybridisation and
the new atomic orbitals are called hybrid orbitals Two key
points about the number and type of hybrid orbitals are that
i The number of hybrid orbitals obtained equals the number ofatomic orbitals mixed
ii The type of hybrid orbitals obtained varies with the types of
atomic orbitals mixed
3 2 3 1 T f h b idi ti
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3231 Type of hybridisation
1 sp3 hybridisation ~ When four electron groups surround
the central atom the center atom involved must prepare fourorbitals with equal energies to overlap with the four
surrounding electron groups Valence Bond theory uses
hypothetical hybrid orbitals which are atomic orbitals
obtained when two or more non-equivalent orbitals of the
same atom combine in preparation for covalent bond
formation H bridisation is the term a lied to the mixin of
atomic orbitals in an atom (usually a central atom) togenerate a set of hybrid orbitals In the case of sp3 We can
generate four equivalent hybrid orbitals from the center
atom by mixing the s orbital and the three p orbitals
Explanation Energy level diagram Diagram of orbitals
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Explanation Energy level diagram Diagram of orbitals
a) Ground state
Carbon which act as thecenter atom has the valence
electron of 2s22p2
b) Excited state One of the electron from 2s
is promote to 2p orbitals -
equa energy eve
c) Hybridised state
One orbital of 2s and three
orbitals of 2p combined(hybrid) and rearrange
themselves to the shape
and orientation of a
tetrahedral shape
Molecular shape Tetrahedral
Angle between bond pair 10950
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2 sp2 hybridisation ~ When three electron groups surround
the central atom the center atom involved must prepare three
orbitals with equal energies to overlap with the three
surrounding electron groups In sp2 hybridisation three
equivalent (in terms of energy level) from the center atom by
mixing the one s orbital and the two p orbitals Using boron
trifluoride (BF3) as example sp2
hybridisation is explained
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8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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3 sp hybridisation ~ When two electron groups surround
the central atom we observe a linear shape which
means that the bonding orbitals must have a linear
orientation VB theory explains this by proposing thatmixing two nonequivalent orbitals of a central atom one
s and one p gives rise to two equivalent sp hybrid
orbitals that lie 1800 apart
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Beryllium which act as the
center atom has the valence
electron of 2s2
b) Excited state
One of the electron from 2s ispromote to 2p orbitals - equal
energy level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of a
linear shape
Molecular shape Linear
Angle between bond pair 1800
4 The concept of hybridisation is also useful to explain molecules with
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p y p
doubletriple bonds By using the concept of the direct overlapping
orbitals and side-touch lapping orbitals the formation of multiple
bonds in ethene C2H4 and ethyne C2H2 are described
a) Ethene C2H4 - Hybridisation take place for both carbon atoms in
ethene molecule is sp2 hybridisation
bull When both hybridised C is bonded together one of the hybridisedorbital overlapped directly between each other while the other two
hybridised orbitals overlapped directly with hydrogen atoms
z
atoms form a side-touch bond between each other and form anotherbond as shown in the diagram below
bull From the diagram with C=C there are two types of bond A sigma-
bond (σ-bond) is covalent bonds formed by orbitals overlapping end-
to-end with the electron density concentrated between the nuclei ofthe bonding atoms while the second type is called a pi bond (π-
bond) which is defined as a covalent bond formed by sideways
overlapping orbitals with electron density concentrated above and
below the plane of the nuclei of the bonding atoms
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital of 2s and two
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of atrigonal planar shape
Note that on unhybridised pz
orbital an electron is presence
Molecular shape Trigonal planar Angle between bond pair 1200
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b) Ethyne C H has the Lewis structure of H CequivC H which the
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b) Ethyne C2H2 has the Lewis structure of HminusCequivCminusH which the
bonding can be explain using sp hybridisation Table below
described how the hybridisation take place on each carbonatom and how the formation of triple bond occur
c) From the diagram formation of -CequivC- is due to the formation
of one sigma-bond by direct overlapping of one of the twohybrid orbitals on each C while the other two pi-bond bonds
are formed as a result of side-lapping of the each two
y z
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital from 2s and 2p
orbitals combined (hybrid) and
rearrange themselves to the
shape and orientation of a
linear shapeNote that on unhybridised py amp
pz orbitals electrons are
presence to form two
π-bonds Molecular shape Linear Angle between bond pair 1800
sp
pz py
5 Other examples of applications in valence bond theory includes the
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formation of nitrogen molecule N2 and hydrogen cyanide HCN
molecule
a) Nitrogen gas is Earths most abundant gas as it cover 78 of the
content of our air Nitrogen molecule is an inert gas thanks to its
short covalent bond and also its strong triple bond Therefore
a lot of heats are required to break the chemical bond of nitrogenbefore it can be applied in industries Using valence bond theory
the bonding of nitrogen molecule is explained in the diagram
b) So when the two hybridised nitrogen atom interacting amongeach other they hence formed a linear shape and two pi-bonds
(πminusbonds) are formed as a result of side-lapping of unhybridised
py and pz orbital respectively
N N
ExplanationEnergy level
diagramDiagram of orbitals
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diagram
a) Ground state
Both nitrogen whichact as the center atom
has the valence
electron of 2s22p3
b) Excited state One of the electron
from 2s is promote to 2p
orbitals - e ual ener
level
c) Hybridised state
One orbital of 2s and
one orbitals of 2p
combined (hybrid) and
rearrange themselves
to the shape and
orientation of a linear
shapeMolecular shape Linear
Angle between bond pair 1800
b) Whereas for hydrogen cyanide HCN both carbon atom and
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) y g y
nitrogen atom undergoes sp hybridisation in order to form a
linear structureExplanation
Energy level diagram for
carbon atom C
Energy level diagram for
nitrogen atom N
a) Ground state
Valence electronC 2s22p2
N 2s22p3
b) Excited state
One of the electron from
2s is promote to 2p
orbitals - equal energy
level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined
(hybrid) and form sp
hybridisation
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N
σσσσ
π
σσσσ
π
6 However there are a few limitation on valence bond theory such
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as when explaining the effect angle of bond-pair and bond-pair
electrons when there isare presence of lone pair electron in thecenter atom and also the difference of electronegativity
a) Bonding in ammonia NH3 and water H2O ~ Both NH3 and
H2O undergoes an arrangement similar to sp3 hybridisation
similar to that of C in methane Table below shows thehybridisation occur for nitrogen in ammonia and oxygen in water
Energy level diagram Energy level diagram
for nitrogen atom N for oxygen atom O
a) Ground state
Valence electron
N 2s22p3
O 2s2
2p4
b) Excited state
One of the electron
from 2s is promote to
2p orbitals
c) Hybridised state
One orbital of 2s and
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One orbital of 2s and
three orbitals from 2p
combined (hybrid) andrearrange themselves to
the shape and
orientation similar to that
of tetrahedral shape
Similar to arrangement in
tetrahedral
Similar to arrangement in
tetrahedral
of ammonia NH3 andwater H2O
Shape and angleShape trigonal pyramidal
Angle 1070
Shape bent
Angle 10450
Number of bond pair amp
lone pair electrons
Bond pair electron 3
Lone pair electron 1
Bond pair electron 2
Lone pair electron 2
bull The angle of bond pair - bond pair electrons in ammonia and0 0
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water are 1070 and 10450 respectively which is lesser than in
methane molecule (10950
) This can be explained by the factof the presence of lone-pair electrons in both ammonia and
water Since the lone pair - lone pair electron repulsion is
stronger than lone pair - bond pair electron repulsion than
bond pair - bond pair electron repulsion it is expected that therepulsion between lone pair - bond pair electrons in ammonia
is stronger than bond pair - bond pair electrons repulsion
hence caused the angle to squeeze to a smaller angle As
for water since there is a presence of lone pair - lone pair
electrons repulsion it results the bond pair - bond pair
electron repulsion to be much smaller hence caused the
angle to squeeze to a smaller angle Though valence bondtheory does not actually explain the hybridisation especially
when it relates to the repulsion occur involving lone pair
electron another theories shall be applied to study such
effect All these shall be discussed further in VSEPR theory
b) Effect of electronegativity and bonding angle between bond pair
- bond pair electrons ~
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- bond pair electrons
i As discussed earlier in the bonding of water (H2
O) the
molecular shape and angle is described in the diagram below
Sul hur which is also an element from Grou 16 formed h dro en
sulphide H2S when sulphur react with hydrogen However unlike
water the bonding angles is much smaller compare to water This is
due to the difference of electronegativity and also the bond length
between O and S in molecule Since O is more electronegative
compare to S the O-H bond is pulled closer toward O Furthermorethe bond length of O-H is shorter compare to S-H As a result the
bonding pair-bonding pair electrons repulse greater with each other in
H-O-H and caused the angle become greater
ii Another example is between NH3 and PH3 The orbital
f
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diagram for both NH3 and PH3 are described below Both
nitrogen N and phosphorous P are from the same groupwhich is Group 15
Similar to the above case P in phosphine is less
electronegative than N in ammonia and bond length of N-H
is shorter than P-H As a result H is pull closer to N and
repulsion between bonding of H-N-H is greater compare toH-P-H and caused the angle between H-N-H is greater
compare to H-P-H
324 Valence Shell Electron Pair Repulsion (VSEPR) Theory
1 M l l t i th th di i l t f t
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1 Molecular geometry is the three-dimensional arrangement of atoms
in a molecule A moleculersquos geometry affects its physical andchemical properties such as melting point boiling point density
and the types of reactions it undergoes The basic concept of
VSEPR theory is based on the three general rules below
a) As far as electron-pair repulsion is concerned double bonds andtriple bonds can be treated like single bonds This approximation is
good for qualitative purposes However you should realize that in
ldquo rdquo
because there are two or three bonds between two atoms theelectron density occupies more space
b) If a molecule has two or more resonance structures we can apply
the VSEPR model to any one of them Formal charges are usually
not shown
c) The order of repulsion strength of lone pair and bond pair are
lone-pair amp lone-pair electrons repulsion are the strongest
followed by lone-pair amp bond-pair electrons repulsion while
bond-pair amp bond-pair electrons repulsion is the weakest
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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bull The interaction between sodium atom and fluorine atom
occur where sodium atom (with low ionisation energy)
donates electron to fluorine (with high electron affinity) to formsodium ion Na+ and fluoride ion F- respectively Note that
both ions have achieved octet arrangement of ns2np6 as it is
the most stable form of an ion formedbull The oppositely charged Na+ and F- form a giant ionic crystal
lattice with very high melting point via electrostatic attraction
i Magnesium chloride MgCl2 ii potassium oxide K2O
iii Calcium sulphide CaS iv aluminium oxide Al2O3
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bull According to Coulumbs Law electrostatic energy between
two oppositely charged substance (A and B) is directly
proportional to the charge carried by each ions yet inversely
proportional to the distance between them
bull This relationship helps us predict trends in lattice energy and
explain the effects of ionic size and charge
( ) minus+
minus+
minus
times
prop r r
energylatticeor EnergyticElectrosta
nn
a) Effect of ionic size As we move down a group in theperiodic table the ionic radius increases Therefore the
electrostatic energy between cations and anions decreases
because the inter-ionic distance is greater thus the lattice
energies of their compounds should decrease as well This
prediction is borne out by the alkali-metal halides note the
regular decrease in lattice energy down a group whether we
hold the cation constant (LiF to LiI)
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b) Effect of ionic charge When we compare lithium fluoride
with magnesium oxide we find cations of about equal radii
(Li+ = 76 pm and Mg2+ = 72 pm) and anions of about equalradii (F- = 133 pm and O2- = 140 pm) Thus the only
significant difference is the ionic charge LiF contains the
singly charged Li+ and F- ions whereas MgO contains the
doubly charged Mg2+ and O2- ions The difference in their
lattice energies is
∆H of LiF = - 1050 kJ mol-1
∆Hlattice of MgO = - 3923 kJ mol-1
This nearly fourfold increase in ∆Hlattice reflects the fourfold
increase in the product of the charges (1 x 1 vs 2 x 2) in the
numerator of equation above
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3 Properties of Ionic Compound
a) Melting point - Ionic compound has giant ionic crystal lattice
which are hold by strong electrostatic attraction forces by
repeating of oppositely charged ions
Therefore the melting point of ionic compounds are usuallyvery high as a lot of energies are required to overcome the
electrostatic attraction forces and melted to form free moving
ions Therefore ionic compounds have very high melting
point
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b) Conductivity of electricity - Most ionic compounds do not
conduct electricity (insulator) in the solid state but do
conduct it when melted or when dissolved in water (exceptsome super-ionic conductors such as AgI which have
remarkable conductivity in the solid state) Solid ionic salt
consists of immobilized ions When it melts or dissolveshowever the ions are free to move and carry an electric
current
-
are hard (does not dent) rigid (does not bend) and brittle(cracks without deforming) These properties are due to the
powerful attractive forces that hold the ions in specific
positions throughout the crystal Moving the ions out of
position requires overcoming these forces so the sample
resists denting and bending If enough pressure is applied
ions of like charge are brought next to each other and
repulsive forces crack the ionic solid suddenly
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32 Covalent bondbull Studies of covalent bond was widely developed ever since
Lewis suggested that a chemical bond exist in a hydrogen gas
occur by sharing en electron between two hydrogen atoms
bull Electron pair that connect the 2 hydrogen atoms is called
covalent bond a bond in which two electrons are shared by
wo a oms an e e ec ron pa r a on e ween e wohydrogen atoms is also called as bonding pair electrons
bull In a covalent bond each electron in a shared pair is attracted to
the nuclei of both atoms This attraction holds the two atoms in
H2 together and is responsible for the formation of covalentbonds in other molecules
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2 A Lewis structure is a representation of covalent bonding
in which shared electron pairs are shown either as lines
or as pairs of dots between two atoms and lone pairsare shown as pairs of dots on individual atoms Only
valence electrons are shown in a Lewis structure
a) Consider the fluorine molecule F2 The electronconfiguration of F is 1s22s22p5 The 1s electrons are in
inner shell which is nearest to the nucleus For this reason
the do not artici ate in bond formation Thus each F
atom has seven valence electrons (2s22p5) Thereforeeach fluorine atom needed one electron to achieve octet
configuration (ns2np6)
F 1 s22s22p5 F 1 s22s22p5 1s2 2s2 2p6 1s2 2s2 2p6
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bull Oxygen atom has electronic configuration of 1s22s22p4 To
achieve stable octet configuration (ns2np6) each oxygen
atom need 2 electrons Hence when 2 oxygen atomsinteract they shared two electrons in between each other as
described in diagram below
O 1s2
2s2
2p4
O 1s2
2s2
2p4
1s2
2s2
2p6
1s2
2s2
2p6
bull From the structure of oxygen molecule formed each oxygen
atom shared two electrons from each other to form a double
bond in order to achieve octet configuration among each
oxygen atom
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bull Nitrogen has electronic configuration of 1s22s22p3 and
required 3 electrons to achieve octet configuration (ns2np6)
In this case each nitrogen atom shared 3 electrons fromeach of its atom to form triple bond
N 1 s22s22p3 N 1 s22s22p3 1s2 2s2 2p6 1s2 2s2 2p6
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Water H2O
H 1 s1 O 2 s22p4
Carbon dioxide CO2
C 2 s22p2 O 2 s22p4
Ammonia NH3
H 1 s1 N 2 s22p3Ethene C2H4
H 1 s1 C 2 s22p2
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Hydrogen cyanide HCN
H 1 s1 N 2 s22p3 C 2 s22p2
Ethanoic acid CH3COOH
H 1 s1 O 2 s22p4 C 2 s22p2
Tetrachloromethane CCl4C 2 s22p2 Cl 3s23p5
Ethyne C2H2
H 1 s1 C 2 s22p2
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a) Note that in ethene hydrogen cyanide and ethanoic acid all
the valence electrons are used in bonding there are no lone
pairs on the carbon atoms In fact most of the stablemolecules containing carbon do not have lone pairs on the
carbon atoms
b) Multiple bonds (double bond or triple bond) are shorter than single covalent bonds Bond length is defined as the
distance between the nuclei of two covalently bonded atoms
nitrogen triple bonds are shorter than double bonds whichin turn are shorter than single bonds The shorter multiple
bonds are also more stable than single bonds
c) Covalent bond not only exist in neutral molecule but also insome molecular ions Table below shows a few example of
molecular ions which have covalent bonds in its molecule
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Carbonate ion CO32-
C 2 s22p2 O 2 s22p4
Cyanide ion CN-
N 2 s22p3 C 2 s22p2
u p a e on 4-
S 3 s23p4 O 2 s22p4
ra e on 3-
N 2 s22p3 O 2 s22p4
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321 Exception of Octet Rules
1 From the Lewis structure sketched for sulphate ion SO42- we can
see that the center metal atom (sulphur) has more than 8electrons For this the molecule is described as molecule that can
expanded octet These molecules that have more than 8
electrons are located at Period 3 and below as these center
atoms have empty d-orbital available to expand the number ofelectrons positioned in the center atom Those center atom from
Period 2 such as C N O and F can only allocate a maximum
Example Both phosphorous and nitrogen are elements from Group15 with the valence electron of ns2np3 They can both form NCl3and PCl3 respectively when react with limited amount of chlorine
however under excess chlorine only PCl 5 can be formed but
not NCl 5 Explain the statement bolded
Solution This is due to phosphorous which is from period
3 have empty d-orbital to expand the octet However
nitrogen which is from Period 2 do not have emptyorbital and can only allocate 8 electrons in its shell
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Phosphorous pentachloride PCl5P 3 s23p3 Cl 3s23p5
Sulphur hexafluoride SF6
S 3 s23p4 F 2 s22p5
Bromine pentachloride BrCl5Cl 3s23p5 Br 4s24p5
Xenon tetrafluoride XeF4
Xe 5s25p6 F 2 s22p5
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2 There are also some stable covalent compounds which have
less than 8 electrons in their center atom (incomplete octet)
The center atoms are usually metals with great number ofvalence electrons with small atomic radius such as
beryllium boron and aluminium
BeCl2Be 2s2 Cl 3s23p5 BF3B 2 s22p1 F 2 s22p5 AlCl3Al 3s23p1 F 3 s23p5
These compound possessed stability due to their short bondlength between center atom with surround atom Furthermore
they can form resonance structure between the center atom
and surrounding atoms
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3 There are also some molecules which contain an odd
number of electrons Among the most common compounds
are nitrogen monoxide (NO) and nitrogen dioxide (NO2)
Odd-electron molecules are sometimes called radicals
Nitrogen monoxide NO
N 2 s22p3 O 2 s22p4
Nitrogen dioxide NO2
N 2 s22p3 O 2 s22p4
any ra ca s are g y reac ve e reason s t at t ere
is a tendency for the unpaired electron to form a
covalent bond with an unpaired electron on another
molecule For example when two nitrogen dioxide
molecules collide they form dinitrogen tetroxide
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322 Dative bond
1 As shown in the Lewis structure of nitrate ion one of the N-Obond is drawn as rarr The bond rarr placed is called as dative
bond (also known as coordinative bond) where dative
bond is defined as a covalent bond in which one of the atoms
donate the lone pair electrons available Although theproperties of a coordinate covalent bond do not differ from
those of a normal covalent bond because all electrons are
a e no ma er w a e r source
2 Dative bond is usually applied for these few circumstances
below
a) To assist atom molecule ion that not yet achieved octet
configuration Making use of atom which has lone pair
electrons to those which are lack of electrons
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Sulphur dioxide SO2Sulphur trioxide SO3 Carbon monoxide CO
Ozone molecule O3 Water with hydrogen ion
S
OO
O
Ammonia with hydrogen ion Ammonia with boron trifluoride
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b) Formation of dimer - In order for some compounds which
have incomplete octet to achieve stability they tend to form
dimer or polymer among themselves by using dative bondTwo of the most common examples are aluminium trichloride
and beryllium dichloride
Monomer of AlCl3Dimer of aluminium chloride Al
2
Cl6
Monomer of BeCl2 Dimer of Be2Cl4 Polymer of (BeCl2)n
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c) Formation of coordination compounds - Coordination
compounds are substances that contain at least one complex
ion a species consisting of a central metal cation (either atransition metal or a main-group metal) that is bonded to
molecules andor anions called ligands via dative
(coordinative) bond
hexaaquacopper (II) ion [Cu(H2O)6]2+ tetraamminenickel (II) ion [Ni(NH3)4]
2+
Hexacyanoferrate (III) ion [Fe(CN)6]3- Trioxalatocobaltate(III)ion [Co(C2O4)3]
3-
3 2 3 Hybridisation Theory
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323 Hybridisation Theory
Valence bond theory
bull The basic principle of valence bond theory is that a covalentbond is formed when orbitals of two atoms overlap and the
overlapped region which is between the nuclei is occupied by
a pair of electrons The central themes of valence bond theory
derive from this principle
ndash Opposing spins of the electron pair ~ Stated in Paulis
orbitals has a maximum capacity of two electrons that musthave opposite spins For example when a covalent bond is
formed in molecule of hydrogen H2 the two 1s electrons of
two H atoms occupy the overlapping 1s orbitals and have
opposite spins
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bull Maximum overlap of bonding orbitals ~ The bond strength
depends on attraction of the nuclei for the shared electrons
so the greater the orbital overlap the stronger (morestable) the bond The extent of overlap depends on the
shapes and directions of the orbitals An s orbital is spherical
but p and d orbitals have more electron density in one
direction than in another Thus whenever possible a bond
involving p or d orbitals will be oriented in the direction that
maximizes overlap For example in hydrogen fluoride (HF)
bond the 1s orbital of H overlaps the half-filled 2p orbital of Falong the long axis of that orbital Any other direction would
result in less overlap and thus a weaker bond
H b idi ti f t i bit l T t f th
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bull Hybridisation of atomic orbitals ~ To account for the
bonding in simple diatomic molecules like HF we picture the
direct overlap of s and p orbitals of isolated atoms But howcan we account for the shapes of so many molecules and
polyatomic ions through the overlap of spherical s orbitals
dumbbell-shaped p orbitals and cloverleaf-shaped d orbitals
Linus Pauling proposed that the valence atomic orbitals in the
molecule are different from those in the isolated atoms The
spatial orientations of these new orbitals lead to more stable
bonds and are consistent with observed molecular shapesThe process of orbital mixing is called hybridisation and
the new atomic orbitals are called hybrid orbitals Two key
points about the number and type of hybrid orbitals are that
i The number of hybrid orbitals obtained equals the number ofatomic orbitals mixed
ii The type of hybrid orbitals obtained varies with the types of
atomic orbitals mixed
3 2 3 1 T f h b idi ti
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3231 Type of hybridisation
1 sp3 hybridisation ~ When four electron groups surround
the central atom the center atom involved must prepare fourorbitals with equal energies to overlap with the four
surrounding electron groups Valence Bond theory uses
hypothetical hybrid orbitals which are atomic orbitals
obtained when two or more non-equivalent orbitals of the
same atom combine in preparation for covalent bond
formation H bridisation is the term a lied to the mixin of
atomic orbitals in an atom (usually a central atom) togenerate a set of hybrid orbitals In the case of sp3 We can
generate four equivalent hybrid orbitals from the center
atom by mixing the s orbital and the three p orbitals
Explanation Energy level diagram Diagram of orbitals
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Explanation Energy level diagram Diagram of orbitals
a) Ground state
Carbon which act as thecenter atom has the valence
electron of 2s22p2
b) Excited state One of the electron from 2s
is promote to 2p orbitals -
equa energy eve
c) Hybridised state
One orbital of 2s and three
orbitals of 2p combined(hybrid) and rearrange
themselves to the shape
and orientation of a
tetrahedral shape
Molecular shape Tetrahedral
Angle between bond pair 10950
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2 sp2 hybridisation ~ When three electron groups surround
the central atom the center atom involved must prepare three
orbitals with equal energies to overlap with the three
surrounding electron groups In sp2 hybridisation three
equivalent (in terms of energy level) from the center atom by
mixing the one s orbital and the two p orbitals Using boron
trifluoride (BF3) as example sp2
hybridisation is explained
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8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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3 sp hybridisation ~ When two electron groups surround
the central atom we observe a linear shape which
means that the bonding orbitals must have a linear
orientation VB theory explains this by proposing thatmixing two nonequivalent orbitals of a central atom one
s and one p gives rise to two equivalent sp hybrid
orbitals that lie 1800 apart
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Beryllium which act as the
center atom has the valence
electron of 2s2
b) Excited state
One of the electron from 2s ispromote to 2p orbitals - equal
energy level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of a
linear shape
Molecular shape Linear
Angle between bond pair 1800
4 The concept of hybridisation is also useful to explain molecules with
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p y p
doubletriple bonds By using the concept of the direct overlapping
orbitals and side-touch lapping orbitals the formation of multiple
bonds in ethene C2H4 and ethyne C2H2 are described
a) Ethene C2H4 - Hybridisation take place for both carbon atoms in
ethene molecule is sp2 hybridisation
bull When both hybridised C is bonded together one of the hybridisedorbital overlapped directly between each other while the other two
hybridised orbitals overlapped directly with hydrogen atoms
z
atoms form a side-touch bond between each other and form anotherbond as shown in the diagram below
bull From the diagram with C=C there are two types of bond A sigma-
bond (σ-bond) is covalent bonds formed by orbitals overlapping end-
to-end with the electron density concentrated between the nuclei ofthe bonding atoms while the second type is called a pi bond (π-
bond) which is defined as a covalent bond formed by sideways
overlapping orbitals with electron density concentrated above and
below the plane of the nuclei of the bonding atoms
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital of 2s and two
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of atrigonal planar shape
Note that on unhybridised pz
orbital an electron is presence
Molecular shape Trigonal planar Angle between bond pair 1200
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b) Ethyne C H has the Lewis structure of H CequivC H which the
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b) Ethyne C2H2 has the Lewis structure of HminusCequivCminusH which the
bonding can be explain using sp hybridisation Table below
described how the hybridisation take place on each carbonatom and how the formation of triple bond occur
c) From the diagram formation of -CequivC- is due to the formation
of one sigma-bond by direct overlapping of one of the twohybrid orbitals on each C while the other two pi-bond bonds
are formed as a result of side-lapping of the each two
y z
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital from 2s and 2p
orbitals combined (hybrid) and
rearrange themselves to the
shape and orientation of a
linear shapeNote that on unhybridised py amp
pz orbitals electrons are
presence to form two
π-bonds Molecular shape Linear Angle between bond pair 1800
sp
pz py
5 Other examples of applications in valence bond theory includes the
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formation of nitrogen molecule N2 and hydrogen cyanide HCN
molecule
a) Nitrogen gas is Earths most abundant gas as it cover 78 of the
content of our air Nitrogen molecule is an inert gas thanks to its
short covalent bond and also its strong triple bond Therefore
a lot of heats are required to break the chemical bond of nitrogenbefore it can be applied in industries Using valence bond theory
the bonding of nitrogen molecule is explained in the diagram
b) So when the two hybridised nitrogen atom interacting amongeach other they hence formed a linear shape and two pi-bonds
(πminusbonds) are formed as a result of side-lapping of unhybridised
py and pz orbital respectively
N N
ExplanationEnergy level
diagramDiagram of orbitals
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diagram
a) Ground state
Both nitrogen whichact as the center atom
has the valence
electron of 2s22p3
b) Excited state One of the electron
from 2s is promote to 2p
orbitals - e ual ener
level
c) Hybridised state
One orbital of 2s and
one orbitals of 2p
combined (hybrid) and
rearrange themselves
to the shape and
orientation of a linear
shapeMolecular shape Linear
Angle between bond pair 1800
b) Whereas for hydrogen cyanide HCN both carbon atom and
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) y g y
nitrogen atom undergoes sp hybridisation in order to form a
linear structureExplanation
Energy level diagram for
carbon atom C
Energy level diagram for
nitrogen atom N
a) Ground state
Valence electronC 2s22p2
N 2s22p3
b) Excited state
One of the electron from
2s is promote to 2p
orbitals - equal energy
level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined
(hybrid) and form sp
hybridisation
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N
σσσσ
π
σσσσ
π
6 However there are a few limitation on valence bond theory such
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as when explaining the effect angle of bond-pair and bond-pair
electrons when there isare presence of lone pair electron in thecenter atom and also the difference of electronegativity
a) Bonding in ammonia NH3 and water H2O ~ Both NH3 and
H2O undergoes an arrangement similar to sp3 hybridisation
similar to that of C in methane Table below shows thehybridisation occur for nitrogen in ammonia and oxygen in water
Energy level diagram Energy level diagram
for nitrogen atom N for oxygen atom O
a) Ground state
Valence electron
N 2s22p3
O 2s2
2p4
b) Excited state
One of the electron
from 2s is promote to
2p orbitals
c) Hybridised state
One orbital of 2s and
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One orbital of 2s and
three orbitals from 2p
combined (hybrid) andrearrange themselves to
the shape and
orientation similar to that
of tetrahedral shape
Similar to arrangement in
tetrahedral
Similar to arrangement in
tetrahedral
of ammonia NH3 andwater H2O
Shape and angleShape trigonal pyramidal
Angle 1070
Shape bent
Angle 10450
Number of bond pair amp
lone pair electrons
Bond pair electron 3
Lone pair electron 1
Bond pair electron 2
Lone pair electron 2
bull The angle of bond pair - bond pair electrons in ammonia and0 0
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water are 1070 and 10450 respectively which is lesser than in
methane molecule (10950
) This can be explained by the factof the presence of lone-pair electrons in both ammonia and
water Since the lone pair - lone pair electron repulsion is
stronger than lone pair - bond pair electron repulsion than
bond pair - bond pair electron repulsion it is expected that therepulsion between lone pair - bond pair electrons in ammonia
is stronger than bond pair - bond pair electrons repulsion
hence caused the angle to squeeze to a smaller angle As
for water since there is a presence of lone pair - lone pair
electrons repulsion it results the bond pair - bond pair
electron repulsion to be much smaller hence caused the
angle to squeeze to a smaller angle Though valence bondtheory does not actually explain the hybridisation especially
when it relates to the repulsion occur involving lone pair
electron another theories shall be applied to study such
effect All these shall be discussed further in VSEPR theory
b) Effect of electronegativity and bonding angle between bond pair
- bond pair electrons ~
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- bond pair electrons
i As discussed earlier in the bonding of water (H2
O) the
molecular shape and angle is described in the diagram below
Sul hur which is also an element from Grou 16 formed h dro en
sulphide H2S when sulphur react with hydrogen However unlike
water the bonding angles is much smaller compare to water This is
due to the difference of electronegativity and also the bond length
between O and S in molecule Since O is more electronegative
compare to S the O-H bond is pulled closer toward O Furthermorethe bond length of O-H is shorter compare to S-H As a result the
bonding pair-bonding pair electrons repulse greater with each other in
H-O-H and caused the angle become greater
ii Another example is between NH3 and PH3 The orbital
f
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diagram for both NH3 and PH3 are described below Both
nitrogen N and phosphorous P are from the same groupwhich is Group 15
Similar to the above case P in phosphine is less
electronegative than N in ammonia and bond length of N-H
is shorter than P-H As a result H is pull closer to N and
repulsion between bonding of H-N-H is greater compare toH-P-H and caused the angle between H-N-H is greater
compare to H-P-H
324 Valence Shell Electron Pair Repulsion (VSEPR) Theory
1 M l l t i th th di i l t f t
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1 Molecular geometry is the three-dimensional arrangement of atoms
in a molecule A moleculersquos geometry affects its physical andchemical properties such as melting point boiling point density
and the types of reactions it undergoes The basic concept of
VSEPR theory is based on the three general rules below
a) As far as electron-pair repulsion is concerned double bonds andtriple bonds can be treated like single bonds This approximation is
good for qualitative purposes However you should realize that in
ldquo rdquo
because there are two or three bonds between two atoms theelectron density occupies more space
b) If a molecule has two or more resonance structures we can apply
the VSEPR model to any one of them Formal charges are usually
not shown
c) The order of repulsion strength of lone pair and bond pair are
lone-pair amp lone-pair electrons repulsion are the strongest
followed by lone-pair amp bond-pair electrons repulsion while
bond-pair amp bond-pair electrons repulsion is the weakest
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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bull According to Coulumbs Law electrostatic energy between
two oppositely charged substance (A and B) is directly
proportional to the charge carried by each ions yet inversely
proportional to the distance between them
bull This relationship helps us predict trends in lattice energy and
explain the effects of ionic size and charge
( ) minus+
minus+
minus
times
prop r r
energylatticeor EnergyticElectrosta
nn
a) Effect of ionic size As we move down a group in theperiodic table the ionic radius increases Therefore the
electrostatic energy between cations and anions decreases
because the inter-ionic distance is greater thus the lattice
energies of their compounds should decrease as well This
prediction is borne out by the alkali-metal halides note the
regular decrease in lattice energy down a group whether we
hold the cation constant (LiF to LiI)
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b) Effect of ionic charge When we compare lithium fluoride
with magnesium oxide we find cations of about equal radii
(Li+ = 76 pm and Mg2+ = 72 pm) and anions of about equalradii (F- = 133 pm and O2- = 140 pm) Thus the only
significant difference is the ionic charge LiF contains the
singly charged Li+ and F- ions whereas MgO contains the
doubly charged Mg2+ and O2- ions The difference in their
lattice energies is
∆H of LiF = - 1050 kJ mol-1
∆Hlattice of MgO = - 3923 kJ mol-1
This nearly fourfold increase in ∆Hlattice reflects the fourfold
increase in the product of the charges (1 x 1 vs 2 x 2) in the
numerator of equation above
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3 Properties of Ionic Compound
a) Melting point - Ionic compound has giant ionic crystal lattice
which are hold by strong electrostatic attraction forces by
repeating of oppositely charged ions
Therefore the melting point of ionic compounds are usuallyvery high as a lot of energies are required to overcome the
electrostatic attraction forces and melted to form free moving
ions Therefore ionic compounds have very high melting
point
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b) Conductivity of electricity - Most ionic compounds do not
conduct electricity (insulator) in the solid state but do
conduct it when melted or when dissolved in water (exceptsome super-ionic conductors such as AgI which have
remarkable conductivity in the solid state) Solid ionic salt
consists of immobilized ions When it melts or dissolveshowever the ions are free to move and carry an electric
current
-
are hard (does not dent) rigid (does not bend) and brittle(cracks without deforming) These properties are due to the
powerful attractive forces that hold the ions in specific
positions throughout the crystal Moving the ions out of
position requires overcoming these forces so the sample
resists denting and bending If enough pressure is applied
ions of like charge are brought next to each other and
repulsive forces crack the ionic solid suddenly
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32 Covalent bondbull Studies of covalent bond was widely developed ever since
Lewis suggested that a chemical bond exist in a hydrogen gas
occur by sharing en electron between two hydrogen atoms
bull Electron pair that connect the 2 hydrogen atoms is called
covalent bond a bond in which two electrons are shared by
wo a oms an e e ec ron pa r a on e ween e wohydrogen atoms is also called as bonding pair electrons
bull In a covalent bond each electron in a shared pair is attracted to
the nuclei of both atoms This attraction holds the two atoms in
H2 together and is responsible for the formation of covalentbonds in other molecules
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2 A Lewis structure is a representation of covalent bonding
in which shared electron pairs are shown either as lines
or as pairs of dots between two atoms and lone pairsare shown as pairs of dots on individual atoms Only
valence electrons are shown in a Lewis structure
a) Consider the fluorine molecule F2 The electronconfiguration of F is 1s22s22p5 The 1s electrons are in
inner shell which is nearest to the nucleus For this reason
the do not artici ate in bond formation Thus each F
atom has seven valence electrons (2s22p5) Thereforeeach fluorine atom needed one electron to achieve octet
configuration (ns2np6)
F 1 s22s22p5 F 1 s22s22p5 1s2 2s2 2p6 1s2 2s2 2p6
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bull Oxygen atom has electronic configuration of 1s22s22p4 To
achieve stable octet configuration (ns2np6) each oxygen
atom need 2 electrons Hence when 2 oxygen atomsinteract they shared two electrons in between each other as
described in diagram below
O 1s2
2s2
2p4
O 1s2
2s2
2p4
1s2
2s2
2p6
1s2
2s2
2p6
bull From the structure of oxygen molecule formed each oxygen
atom shared two electrons from each other to form a double
bond in order to achieve octet configuration among each
oxygen atom
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bull Nitrogen has electronic configuration of 1s22s22p3 and
required 3 electrons to achieve octet configuration (ns2np6)
In this case each nitrogen atom shared 3 electrons fromeach of its atom to form triple bond
N 1 s22s22p3 N 1 s22s22p3 1s2 2s2 2p6 1s2 2s2 2p6
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Water H2O
H 1 s1 O 2 s22p4
Carbon dioxide CO2
C 2 s22p2 O 2 s22p4
Ammonia NH3
H 1 s1 N 2 s22p3Ethene C2H4
H 1 s1 C 2 s22p2
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Hydrogen cyanide HCN
H 1 s1 N 2 s22p3 C 2 s22p2
Ethanoic acid CH3COOH
H 1 s1 O 2 s22p4 C 2 s22p2
Tetrachloromethane CCl4C 2 s22p2 Cl 3s23p5
Ethyne C2H2
H 1 s1 C 2 s22p2
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a) Note that in ethene hydrogen cyanide and ethanoic acid all
the valence electrons are used in bonding there are no lone
pairs on the carbon atoms In fact most of the stablemolecules containing carbon do not have lone pairs on the
carbon atoms
b) Multiple bonds (double bond or triple bond) are shorter than single covalent bonds Bond length is defined as the
distance between the nuclei of two covalently bonded atoms
nitrogen triple bonds are shorter than double bonds whichin turn are shorter than single bonds The shorter multiple
bonds are also more stable than single bonds
c) Covalent bond not only exist in neutral molecule but also insome molecular ions Table below shows a few example of
molecular ions which have covalent bonds in its molecule
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Carbonate ion CO32-
C 2 s22p2 O 2 s22p4
Cyanide ion CN-
N 2 s22p3 C 2 s22p2
u p a e on 4-
S 3 s23p4 O 2 s22p4
ra e on 3-
N 2 s22p3 O 2 s22p4
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321 Exception of Octet Rules
1 From the Lewis structure sketched for sulphate ion SO42- we can
see that the center metal atom (sulphur) has more than 8electrons For this the molecule is described as molecule that can
expanded octet These molecules that have more than 8
electrons are located at Period 3 and below as these center
atoms have empty d-orbital available to expand the number ofelectrons positioned in the center atom Those center atom from
Period 2 such as C N O and F can only allocate a maximum
Example Both phosphorous and nitrogen are elements from Group15 with the valence electron of ns2np3 They can both form NCl3and PCl3 respectively when react with limited amount of chlorine
however under excess chlorine only PCl 5 can be formed but
not NCl 5 Explain the statement bolded
Solution This is due to phosphorous which is from period
3 have empty d-orbital to expand the octet However
nitrogen which is from Period 2 do not have emptyorbital and can only allocate 8 electrons in its shell
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Phosphorous pentachloride PCl5P 3 s23p3 Cl 3s23p5
Sulphur hexafluoride SF6
S 3 s23p4 F 2 s22p5
Bromine pentachloride BrCl5Cl 3s23p5 Br 4s24p5
Xenon tetrafluoride XeF4
Xe 5s25p6 F 2 s22p5
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2 There are also some stable covalent compounds which have
less than 8 electrons in their center atom (incomplete octet)
The center atoms are usually metals with great number ofvalence electrons with small atomic radius such as
beryllium boron and aluminium
BeCl2Be 2s2 Cl 3s23p5 BF3B 2 s22p1 F 2 s22p5 AlCl3Al 3s23p1 F 3 s23p5
These compound possessed stability due to their short bondlength between center atom with surround atom Furthermore
they can form resonance structure between the center atom
and surrounding atoms
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3 There are also some molecules which contain an odd
number of electrons Among the most common compounds
are nitrogen monoxide (NO) and nitrogen dioxide (NO2)
Odd-electron molecules are sometimes called radicals
Nitrogen monoxide NO
N 2 s22p3 O 2 s22p4
Nitrogen dioxide NO2
N 2 s22p3 O 2 s22p4
any ra ca s are g y reac ve e reason s t at t ere
is a tendency for the unpaired electron to form a
covalent bond with an unpaired electron on another
molecule For example when two nitrogen dioxide
molecules collide they form dinitrogen tetroxide
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322 Dative bond
1 As shown in the Lewis structure of nitrate ion one of the N-Obond is drawn as rarr The bond rarr placed is called as dative
bond (also known as coordinative bond) where dative
bond is defined as a covalent bond in which one of the atoms
donate the lone pair electrons available Although theproperties of a coordinate covalent bond do not differ from
those of a normal covalent bond because all electrons are
a e no ma er w a e r source
2 Dative bond is usually applied for these few circumstances
below
a) To assist atom molecule ion that not yet achieved octet
configuration Making use of atom which has lone pair
electrons to those which are lack of electrons
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Sulphur dioxide SO2Sulphur trioxide SO3 Carbon monoxide CO
Ozone molecule O3 Water with hydrogen ion
S
OO
O
Ammonia with hydrogen ion Ammonia with boron trifluoride
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b) Formation of dimer - In order for some compounds which
have incomplete octet to achieve stability they tend to form
dimer or polymer among themselves by using dative bondTwo of the most common examples are aluminium trichloride
and beryllium dichloride
Monomer of AlCl3Dimer of aluminium chloride Al
2
Cl6
Monomer of BeCl2 Dimer of Be2Cl4 Polymer of (BeCl2)n
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c) Formation of coordination compounds - Coordination
compounds are substances that contain at least one complex
ion a species consisting of a central metal cation (either atransition metal or a main-group metal) that is bonded to
molecules andor anions called ligands via dative
(coordinative) bond
hexaaquacopper (II) ion [Cu(H2O)6]2+ tetraamminenickel (II) ion [Ni(NH3)4]
2+
Hexacyanoferrate (III) ion [Fe(CN)6]3- Trioxalatocobaltate(III)ion [Co(C2O4)3]
3-
3 2 3 Hybridisation Theory
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323 Hybridisation Theory
Valence bond theory
bull The basic principle of valence bond theory is that a covalentbond is formed when orbitals of two atoms overlap and the
overlapped region which is between the nuclei is occupied by
a pair of electrons The central themes of valence bond theory
derive from this principle
ndash Opposing spins of the electron pair ~ Stated in Paulis
orbitals has a maximum capacity of two electrons that musthave opposite spins For example when a covalent bond is
formed in molecule of hydrogen H2 the two 1s electrons of
two H atoms occupy the overlapping 1s orbitals and have
opposite spins
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bull Maximum overlap of bonding orbitals ~ The bond strength
depends on attraction of the nuclei for the shared electrons
so the greater the orbital overlap the stronger (morestable) the bond The extent of overlap depends on the
shapes and directions of the orbitals An s orbital is spherical
but p and d orbitals have more electron density in one
direction than in another Thus whenever possible a bond
involving p or d orbitals will be oriented in the direction that
maximizes overlap For example in hydrogen fluoride (HF)
bond the 1s orbital of H overlaps the half-filled 2p orbital of Falong the long axis of that orbital Any other direction would
result in less overlap and thus a weaker bond
H b idi ti f t i bit l T t f th
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bull Hybridisation of atomic orbitals ~ To account for the
bonding in simple diatomic molecules like HF we picture the
direct overlap of s and p orbitals of isolated atoms But howcan we account for the shapes of so many molecules and
polyatomic ions through the overlap of spherical s orbitals
dumbbell-shaped p orbitals and cloverleaf-shaped d orbitals
Linus Pauling proposed that the valence atomic orbitals in the
molecule are different from those in the isolated atoms The
spatial orientations of these new orbitals lead to more stable
bonds and are consistent with observed molecular shapesThe process of orbital mixing is called hybridisation and
the new atomic orbitals are called hybrid orbitals Two key
points about the number and type of hybrid orbitals are that
i The number of hybrid orbitals obtained equals the number ofatomic orbitals mixed
ii The type of hybrid orbitals obtained varies with the types of
atomic orbitals mixed
3 2 3 1 T f h b idi ti
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3231 Type of hybridisation
1 sp3 hybridisation ~ When four electron groups surround
the central atom the center atom involved must prepare fourorbitals with equal energies to overlap with the four
surrounding electron groups Valence Bond theory uses
hypothetical hybrid orbitals which are atomic orbitals
obtained when two or more non-equivalent orbitals of the
same atom combine in preparation for covalent bond
formation H bridisation is the term a lied to the mixin of
atomic orbitals in an atom (usually a central atom) togenerate a set of hybrid orbitals In the case of sp3 We can
generate four equivalent hybrid orbitals from the center
atom by mixing the s orbital and the three p orbitals
Explanation Energy level diagram Diagram of orbitals
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Explanation Energy level diagram Diagram of orbitals
a) Ground state
Carbon which act as thecenter atom has the valence
electron of 2s22p2
b) Excited state One of the electron from 2s
is promote to 2p orbitals -
equa energy eve
c) Hybridised state
One orbital of 2s and three
orbitals of 2p combined(hybrid) and rearrange
themselves to the shape
and orientation of a
tetrahedral shape
Molecular shape Tetrahedral
Angle between bond pair 10950
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2 sp2 hybridisation ~ When three electron groups surround
the central atom the center atom involved must prepare three
orbitals with equal energies to overlap with the three
surrounding electron groups In sp2 hybridisation three
equivalent (in terms of energy level) from the center atom by
mixing the one s orbital and the two p orbitals Using boron
trifluoride (BF3) as example sp2
hybridisation is explained
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8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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3 sp hybridisation ~ When two electron groups surround
the central atom we observe a linear shape which
means that the bonding orbitals must have a linear
orientation VB theory explains this by proposing thatmixing two nonequivalent orbitals of a central atom one
s and one p gives rise to two equivalent sp hybrid
orbitals that lie 1800 apart
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Beryllium which act as the
center atom has the valence
electron of 2s2
b) Excited state
One of the electron from 2s ispromote to 2p orbitals - equal
energy level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of a
linear shape
Molecular shape Linear
Angle between bond pair 1800
4 The concept of hybridisation is also useful to explain molecules with
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p y p
doubletriple bonds By using the concept of the direct overlapping
orbitals and side-touch lapping orbitals the formation of multiple
bonds in ethene C2H4 and ethyne C2H2 are described
a) Ethene C2H4 - Hybridisation take place for both carbon atoms in
ethene molecule is sp2 hybridisation
bull When both hybridised C is bonded together one of the hybridisedorbital overlapped directly between each other while the other two
hybridised orbitals overlapped directly with hydrogen atoms
z
atoms form a side-touch bond between each other and form anotherbond as shown in the diagram below
bull From the diagram with C=C there are two types of bond A sigma-
bond (σ-bond) is covalent bonds formed by orbitals overlapping end-
to-end with the electron density concentrated between the nuclei ofthe bonding atoms while the second type is called a pi bond (π-
bond) which is defined as a covalent bond formed by sideways
overlapping orbitals with electron density concentrated above and
below the plane of the nuclei of the bonding atoms
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital of 2s and two
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of atrigonal planar shape
Note that on unhybridised pz
orbital an electron is presence
Molecular shape Trigonal planar Angle between bond pair 1200
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b) Ethyne C H has the Lewis structure of H CequivC H which the
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b) Ethyne C2H2 has the Lewis structure of HminusCequivCminusH which the
bonding can be explain using sp hybridisation Table below
described how the hybridisation take place on each carbonatom and how the formation of triple bond occur
c) From the diagram formation of -CequivC- is due to the formation
of one sigma-bond by direct overlapping of one of the twohybrid orbitals on each C while the other two pi-bond bonds
are formed as a result of side-lapping of the each two
y z
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital from 2s and 2p
orbitals combined (hybrid) and
rearrange themselves to the
shape and orientation of a
linear shapeNote that on unhybridised py amp
pz orbitals electrons are
presence to form two
π-bonds Molecular shape Linear Angle between bond pair 1800
sp
pz py
5 Other examples of applications in valence bond theory includes the
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formation of nitrogen molecule N2 and hydrogen cyanide HCN
molecule
a) Nitrogen gas is Earths most abundant gas as it cover 78 of the
content of our air Nitrogen molecule is an inert gas thanks to its
short covalent bond and also its strong triple bond Therefore
a lot of heats are required to break the chemical bond of nitrogenbefore it can be applied in industries Using valence bond theory
the bonding of nitrogen molecule is explained in the diagram
b) So when the two hybridised nitrogen atom interacting amongeach other they hence formed a linear shape and two pi-bonds
(πminusbonds) are formed as a result of side-lapping of unhybridised
py and pz orbital respectively
N N
ExplanationEnergy level
diagramDiagram of orbitals
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diagram
a) Ground state
Both nitrogen whichact as the center atom
has the valence
electron of 2s22p3
b) Excited state One of the electron
from 2s is promote to 2p
orbitals - e ual ener
level
c) Hybridised state
One orbital of 2s and
one orbitals of 2p
combined (hybrid) and
rearrange themselves
to the shape and
orientation of a linear
shapeMolecular shape Linear
Angle between bond pair 1800
b) Whereas for hydrogen cyanide HCN both carbon atom and
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) y g y
nitrogen atom undergoes sp hybridisation in order to form a
linear structureExplanation
Energy level diagram for
carbon atom C
Energy level diagram for
nitrogen atom N
a) Ground state
Valence electronC 2s22p2
N 2s22p3
b) Excited state
One of the electron from
2s is promote to 2p
orbitals - equal energy
level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined
(hybrid) and form sp
hybridisation
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N
σσσσ
π
σσσσ
π
6 However there are a few limitation on valence bond theory such
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as when explaining the effect angle of bond-pair and bond-pair
electrons when there isare presence of lone pair electron in thecenter atom and also the difference of electronegativity
a) Bonding in ammonia NH3 and water H2O ~ Both NH3 and
H2O undergoes an arrangement similar to sp3 hybridisation
similar to that of C in methane Table below shows thehybridisation occur for nitrogen in ammonia and oxygen in water
Energy level diagram Energy level diagram
for nitrogen atom N for oxygen atom O
a) Ground state
Valence electron
N 2s22p3
O 2s2
2p4
b) Excited state
One of the electron
from 2s is promote to
2p orbitals
c) Hybridised state
One orbital of 2s and
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One orbital of 2s and
three orbitals from 2p
combined (hybrid) andrearrange themselves to
the shape and
orientation similar to that
of tetrahedral shape
Similar to arrangement in
tetrahedral
Similar to arrangement in
tetrahedral
of ammonia NH3 andwater H2O
Shape and angleShape trigonal pyramidal
Angle 1070
Shape bent
Angle 10450
Number of bond pair amp
lone pair electrons
Bond pair electron 3
Lone pair electron 1
Bond pair electron 2
Lone pair electron 2
bull The angle of bond pair - bond pair electrons in ammonia and0 0
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water are 1070 and 10450 respectively which is lesser than in
methane molecule (10950
) This can be explained by the factof the presence of lone-pair electrons in both ammonia and
water Since the lone pair - lone pair electron repulsion is
stronger than lone pair - bond pair electron repulsion than
bond pair - bond pair electron repulsion it is expected that therepulsion between lone pair - bond pair electrons in ammonia
is stronger than bond pair - bond pair electrons repulsion
hence caused the angle to squeeze to a smaller angle As
for water since there is a presence of lone pair - lone pair
electrons repulsion it results the bond pair - bond pair
electron repulsion to be much smaller hence caused the
angle to squeeze to a smaller angle Though valence bondtheory does not actually explain the hybridisation especially
when it relates to the repulsion occur involving lone pair
electron another theories shall be applied to study such
effect All these shall be discussed further in VSEPR theory
b) Effect of electronegativity and bonding angle between bond pair
- bond pair electrons ~
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- bond pair electrons
i As discussed earlier in the bonding of water (H2
O) the
molecular shape and angle is described in the diagram below
Sul hur which is also an element from Grou 16 formed h dro en
sulphide H2S when sulphur react with hydrogen However unlike
water the bonding angles is much smaller compare to water This is
due to the difference of electronegativity and also the bond length
between O and S in molecule Since O is more electronegative
compare to S the O-H bond is pulled closer toward O Furthermorethe bond length of O-H is shorter compare to S-H As a result the
bonding pair-bonding pair electrons repulse greater with each other in
H-O-H and caused the angle become greater
ii Another example is between NH3 and PH3 The orbital
f
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diagram for both NH3 and PH3 are described below Both
nitrogen N and phosphorous P are from the same groupwhich is Group 15
Similar to the above case P in phosphine is less
electronegative than N in ammonia and bond length of N-H
is shorter than P-H As a result H is pull closer to N and
repulsion between bonding of H-N-H is greater compare toH-P-H and caused the angle between H-N-H is greater
compare to H-P-H
324 Valence Shell Electron Pair Repulsion (VSEPR) Theory
1 M l l t i th th di i l t f t
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1 Molecular geometry is the three-dimensional arrangement of atoms
in a molecule A moleculersquos geometry affects its physical andchemical properties such as melting point boiling point density
and the types of reactions it undergoes The basic concept of
VSEPR theory is based on the three general rules below
a) As far as electron-pair repulsion is concerned double bonds andtriple bonds can be treated like single bonds This approximation is
good for qualitative purposes However you should realize that in
ldquo rdquo
because there are two or three bonds between two atoms theelectron density occupies more space
b) If a molecule has two or more resonance structures we can apply
the VSEPR model to any one of them Formal charges are usually
not shown
c) The order of repulsion strength of lone pair and bond pair are
lone-pair amp lone-pair electrons repulsion are the strongest
followed by lone-pair amp bond-pair electrons repulsion while
bond-pair amp bond-pair electrons repulsion is the weakest
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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b) Effect of ionic charge When we compare lithium fluoride
with magnesium oxide we find cations of about equal radii
(Li+ = 76 pm and Mg2+ = 72 pm) and anions of about equalradii (F- = 133 pm and O2- = 140 pm) Thus the only
significant difference is the ionic charge LiF contains the
singly charged Li+ and F- ions whereas MgO contains the
doubly charged Mg2+ and O2- ions The difference in their
lattice energies is
∆H of LiF = - 1050 kJ mol-1
∆Hlattice of MgO = - 3923 kJ mol-1
This nearly fourfold increase in ∆Hlattice reflects the fourfold
increase in the product of the charges (1 x 1 vs 2 x 2) in the
numerator of equation above
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3 Properties of Ionic Compound
a) Melting point - Ionic compound has giant ionic crystal lattice
which are hold by strong electrostatic attraction forces by
repeating of oppositely charged ions
Therefore the melting point of ionic compounds are usuallyvery high as a lot of energies are required to overcome the
electrostatic attraction forces and melted to form free moving
ions Therefore ionic compounds have very high melting
point
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b) Conductivity of electricity - Most ionic compounds do not
conduct electricity (insulator) in the solid state but do
conduct it when melted or when dissolved in water (exceptsome super-ionic conductors such as AgI which have
remarkable conductivity in the solid state) Solid ionic salt
consists of immobilized ions When it melts or dissolveshowever the ions are free to move and carry an electric
current
-
are hard (does not dent) rigid (does not bend) and brittle(cracks without deforming) These properties are due to the
powerful attractive forces that hold the ions in specific
positions throughout the crystal Moving the ions out of
position requires overcoming these forces so the sample
resists denting and bending If enough pressure is applied
ions of like charge are brought next to each other and
repulsive forces crack the ionic solid suddenly
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32 Covalent bondbull Studies of covalent bond was widely developed ever since
Lewis suggested that a chemical bond exist in a hydrogen gas
occur by sharing en electron between two hydrogen atoms
bull Electron pair that connect the 2 hydrogen atoms is called
covalent bond a bond in which two electrons are shared by
wo a oms an e e ec ron pa r a on e ween e wohydrogen atoms is also called as bonding pair electrons
bull In a covalent bond each electron in a shared pair is attracted to
the nuclei of both atoms This attraction holds the two atoms in
H2 together and is responsible for the formation of covalentbonds in other molecules
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2 A Lewis structure is a representation of covalent bonding
in which shared electron pairs are shown either as lines
or as pairs of dots between two atoms and lone pairsare shown as pairs of dots on individual atoms Only
valence electrons are shown in a Lewis structure
a) Consider the fluorine molecule F2 The electronconfiguration of F is 1s22s22p5 The 1s electrons are in
inner shell which is nearest to the nucleus For this reason
the do not artici ate in bond formation Thus each F
atom has seven valence electrons (2s22p5) Thereforeeach fluorine atom needed one electron to achieve octet
configuration (ns2np6)
F 1 s22s22p5 F 1 s22s22p5 1s2 2s2 2p6 1s2 2s2 2p6
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bull Oxygen atom has electronic configuration of 1s22s22p4 To
achieve stable octet configuration (ns2np6) each oxygen
atom need 2 electrons Hence when 2 oxygen atomsinteract they shared two electrons in between each other as
described in diagram below
O 1s2
2s2
2p4
O 1s2
2s2
2p4
1s2
2s2
2p6
1s2
2s2
2p6
bull From the structure of oxygen molecule formed each oxygen
atom shared two electrons from each other to form a double
bond in order to achieve octet configuration among each
oxygen atom
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bull Nitrogen has electronic configuration of 1s22s22p3 and
required 3 electrons to achieve octet configuration (ns2np6)
In this case each nitrogen atom shared 3 electrons fromeach of its atom to form triple bond
N 1 s22s22p3 N 1 s22s22p3 1s2 2s2 2p6 1s2 2s2 2p6
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Water H2O
H 1 s1 O 2 s22p4
Carbon dioxide CO2
C 2 s22p2 O 2 s22p4
Ammonia NH3
H 1 s1 N 2 s22p3Ethene C2H4
H 1 s1 C 2 s22p2
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Hydrogen cyanide HCN
H 1 s1 N 2 s22p3 C 2 s22p2
Ethanoic acid CH3COOH
H 1 s1 O 2 s22p4 C 2 s22p2
Tetrachloromethane CCl4C 2 s22p2 Cl 3s23p5
Ethyne C2H2
H 1 s1 C 2 s22p2
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a) Note that in ethene hydrogen cyanide and ethanoic acid all
the valence electrons are used in bonding there are no lone
pairs on the carbon atoms In fact most of the stablemolecules containing carbon do not have lone pairs on the
carbon atoms
b) Multiple bonds (double bond or triple bond) are shorter than single covalent bonds Bond length is defined as the
distance between the nuclei of two covalently bonded atoms
nitrogen triple bonds are shorter than double bonds whichin turn are shorter than single bonds The shorter multiple
bonds are also more stable than single bonds
c) Covalent bond not only exist in neutral molecule but also insome molecular ions Table below shows a few example of
molecular ions which have covalent bonds in its molecule
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Carbonate ion CO32-
C 2 s22p2 O 2 s22p4
Cyanide ion CN-
N 2 s22p3 C 2 s22p2
u p a e on 4-
S 3 s23p4 O 2 s22p4
ra e on 3-
N 2 s22p3 O 2 s22p4
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321 Exception of Octet Rules
1 From the Lewis structure sketched for sulphate ion SO42- we can
see that the center metal atom (sulphur) has more than 8electrons For this the molecule is described as molecule that can
expanded octet These molecules that have more than 8
electrons are located at Period 3 and below as these center
atoms have empty d-orbital available to expand the number ofelectrons positioned in the center atom Those center atom from
Period 2 such as C N O and F can only allocate a maximum
Example Both phosphorous and nitrogen are elements from Group15 with the valence electron of ns2np3 They can both form NCl3and PCl3 respectively when react with limited amount of chlorine
however under excess chlorine only PCl 5 can be formed but
not NCl 5 Explain the statement bolded
Solution This is due to phosphorous which is from period
3 have empty d-orbital to expand the octet However
nitrogen which is from Period 2 do not have emptyorbital and can only allocate 8 electrons in its shell
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Phosphorous pentachloride PCl5P 3 s23p3 Cl 3s23p5
Sulphur hexafluoride SF6
S 3 s23p4 F 2 s22p5
Bromine pentachloride BrCl5Cl 3s23p5 Br 4s24p5
Xenon tetrafluoride XeF4
Xe 5s25p6 F 2 s22p5
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2 There are also some stable covalent compounds which have
less than 8 electrons in their center atom (incomplete octet)
The center atoms are usually metals with great number ofvalence electrons with small atomic radius such as
beryllium boron and aluminium
BeCl2Be 2s2 Cl 3s23p5 BF3B 2 s22p1 F 2 s22p5 AlCl3Al 3s23p1 F 3 s23p5
These compound possessed stability due to their short bondlength between center atom with surround atom Furthermore
they can form resonance structure between the center atom
and surrounding atoms
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3 There are also some molecules which contain an odd
number of electrons Among the most common compounds
are nitrogen monoxide (NO) and nitrogen dioxide (NO2)
Odd-electron molecules are sometimes called radicals
Nitrogen monoxide NO
N 2 s22p3 O 2 s22p4
Nitrogen dioxide NO2
N 2 s22p3 O 2 s22p4
any ra ca s are g y reac ve e reason s t at t ere
is a tendency for the unpaired electron to form a
covalent bond with an unpaired electron on another
molecule For example when two nitrogen dioxide
molecules collide they form dinitrogen tetroxide
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322 Dative bond
1 As shown in the Lewis structure of nitrate ion one of the N-Obond is drawn as rarr The bond rarr placed is called as dative
bond (also known as coordinative bond) where dative
bond is defined as a covalent bond in which one of the atoms
donate the lone pair electrons available Although theproperties of a coordinate covalent bond do not differ from
those of a normal covalent bond because all electrons are
a e no ma er w a e r source
2 Dative bond is usually applied for these few circumstances
below
a) To assist atom molecule ion that not yet achieved octet
configuration Making use of atom which has lone pair
electrons to those which are lack of electrons
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Sulphur dioxide SO2Sulphur trioxide SO3 Carbon monoxide CO
Ozone molecule O3 Water with hydrogen ion
S
OO
O
Ammonia with hydrogen ion Ammonia with boron trifluoride
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b) Formation of dimer - In order for some compounds which
have incomplete octet to achieve stability they tend to form
dimer or polymer among themselves by using dative bondTwo of the most common examples are aluminium trichloride
and beryllium dichloride
Monomer of AlCl3Dimer of aluminium chloride Al
2
Cl6
Monomer of BeCl2 Dimer of Be2Cl4 Polymer of (BeCl2)n
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c) Formation of coordination compounds - Coordination
compounds are substances that contain at least one complex
ion a species consisting of a central metal cation (either atransition metal or a main-group metal) that is bonded to
molecules andor anions called ligands via dative
(coordinative) bond
hexaaquacopper (II) ion [Cu(H2O)6]2+ tetraamminenickel (II) ion [Ni(NH3)4]
2+
Hexacyanoferrate (III) ion [Fe(CN)6]3- Trioxalatocobaltate(III)ion [Co(C2O4)3]
3-
3 2 3 Hybridisation Theory
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323 Hybridisation Theory
Valence bond theory
bull The basic principle of valence bond theory is that a covalentbond is formed when orbitals of two atoms overlap and the
overlapped region which is between the nuclei is occupied by
a pair of electrons The central themes of valence bond theory
derive from this principle
ndash Opposing spins of the electron pair ~ Stated in Paulis
orbitals has a maximum capacity of two electrons that musthave opposite spins For example when a covalent bond is
formed in molecule of hydrogen H2 the two 1s electrons of
two H atoms occupy the overlapping 1s orbitals and have
opposite spins
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bull Maximum overlap of bonding orbitals ~ The bond strength
depends on attraction of the nuclei for the shared electrons
so the greater the orbital overlap the stronger (morestable) the bond The extent of overlap depends on the
shapes and directions of the orbitals An s orbital is spherical
but p and d orbitals have more electron density in one
direction than in another Thus whenever possible a bond
involving p or d orbitals will be oriented in the direction that
maximizes overlap For example in hydrogen fluoride (HF)
bond the 1s orbital of H overlaps the half-filled 2p orbital of Falong the long axis of that orbital Any other direction would
result in less overlap and thus a weaker bond
H b idi ti f t i bit l T t f th
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bull Hybridisation of atomic orbitals ~ To account for the
bonding in simple diatomic molecules like HF we picture the
direct overlap of s and p orbitals of isolated atoms But howcan we account for the shapes of so many molecules and
polyatomic ions through the overlap of spherical s orbitals
dumbbell-shaped p orbitals and cloverleaf-shaped d orbitals
Linus Pauling proposed that the valence atomic orbitals in the
molecule are different from those in the isolated atoms The
spatial orientations of these new orbitals lead to more stable
bonds and are consistent with observed molecular shapesThe process of orbital mixing is called hybridisation and
the new atomic orbitals are called hybrid orbitals Two key
points about the number and type of hybrid orbitals are that
i The number of hybrid orbitals obtained equals the number ofatomic orbitals mixed
ii The type of hybrid orbitals obtained varies with the types of
atomic orbitals mixed
3 2 3 1 T f h b idi ti
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3231 Type of hybridisation
1 sp3 hybridisation ~ When four electron groups surround
the central atom the center atom involved must prepare fourorbitals with equal energies to overlap with the four
surrounding electron groups Valence Bond theory uses
hypothetical hybrid orbitals which are atomic orbitals
obtained when two or more non-equivalent orbitals of the
same atom combine in preparation for covalent bond
formation H bridisation is the term a lied to the mixin of
atomic orbitals in an atom (usually a central atom) togenerate a set of hybrid orbitals In the case of sp3 We can
generate four equivalent hybrid orbitals from the center
atom by mixing the s orbital and the three p orbitals
Explanation Energy level diagram Diagram of orbitals
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Explanation Energy level diagram Diagram of orbitals
a) Ground state
Carbon which act as thecenter atom has the valence
electron of 2s22p2
b) Excited state One of the electron from 2s
is promote to 2p orbitals -
equa energy eve
c) Hybridised state
One orbital of 2s and three
orbitals of 2p combined(hybrid) and rearrange
themselves to the shape
and orientation of a
tetrahedral shape
Molecular shape Tetrahedral
Angle between bond pair 10950
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2 sp2 hybridisation ~ When three electron groups surround
the central atom the center atom involved must prepare three
orbitals with equal energies to overlap with the three
surrounding electron groups In sp2 hybridisation three
equivalent (in terms of energy level) from the center atom by
mixing the one s orbital and the two p orbitals Using boron
trifluoride (BF3) as example sp2
hybridisation is explained
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8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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3 sp hybridisation ~ When two electron groups surround
the central atom we observe a linear shape which
means that the bonding orbitals must have a linear
orientation VB theory explains this by proposing thatmixing two nonequivalent orbitals of a central atom one
s and one p gives rise to two equivalent sp hybrid
orbitals that lie 1800 apart
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Beryllium which act as the
center atom has the valence
electron of 2s2
b) Excited state
One of the electron from 2s ispromote to 2p orbitals - equal
energy level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of a
linear shape
Molecular shape Linear
Angle between bond pair 1800
4 The concept of hybridisation is also useful to explain molecules with
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p y p
doubletriple bonds By using the concept of the direct overlapping
orbitals and side-touch lapping orbitals the formation of multiple
bonds in ethene C2H4 and ethyne C2H2 are described
a) Ethene C2H4 - Hybridisation take place for both carbon atoms in
ethene molecule is sp2 hybridisation
bull When both hybridised C is bonded together one of the hybridisedorbital overlapped directly between each other while the other two
hybridised orbitals overlapped directly with hydrogen atoms
z
atoms form a side-touch bond between each other and form anotherbond as shown in the diagram below
bull From the diagram with C=C there are two types of bond A sigma-
bond (σ-bond) is covalent bonds formed by orbitals overlapping end-
to-end with the electron density concentrated between the nuclei ofthe bonding atoms while the second type is called a pi bond (π-
bond) which is defined as a covalent bond formed by sideways
overlapping orbitals with electron density concentrated above and
below the plane of the nuclei of the bonding atoms
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital of 2s and two
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of atrigonal planar shape
Note that on unhybridised pz
orbital an electron is presence
Molecular shape Trigonal planar Angle between bond pair 1200
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b) Ethyne C H has the Lewis structure of H CequivC H which the
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b) Ethyne C2H2 has the Lewis structure of HminusCequivCminusH which the
bonding can be explain using sp hybridisation Table below
described how the hybridisation take place on each carbonatom and how the formation of triple bond occur
c) From the diagram formation of -CequivC- is due to the formation
of one sigma-bond by direct overlapping of one of the twohybrid orbitals on each C while the other two pi-bond bonds
are formed as a result of side-lapping of the each two
y z
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital from 2s and 2p
orbitals combined (hybrid) and
rearrange themselves to the
shape and orientation of a
linear shapeNote that on unhybridised py amp
pz orbitals electrons are
presence to form two
π-bonds Molecular shape Linear Angle between bond pair 1800
sp
pz py
5 Other examples of applications in valence bond theory includes the
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formation of nitrogen molecule N2 and hydrogen cyanide HCN
molecule
a) Nitrogen gas is Earths most abundant gas as it cover 78 of the
content of our air Nitrogen molecule is an inert gas thanks to its
short covalent bond and also its strong triple bond Therefore
a lot of heats are required to break the chemical bond of nitrogenbefore it can be applied in industries Using valence bond theory
the bonding of nitrogen molecule is explained in the diagram
b) So when the two hybridised nitrogen atom interacting amongeach other they hence formed a linear shape and two pi-bonds
(πminusbonds) are formed as a result of side-lapping of unhybridised
py and pz orbital respectively
N N
ExplanationEnergy level
diagramDiagram of orbitals
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diagram
a) Ground state
Both nitrogen whichact as the center atom
has the valence
electron of 2s22p3
b) Excited state One of the electron
from 2s is promote to 2p
orbitals - e ual ener
level
c) Hybridised state
One orbital of 2s and
one orbitals of 2p
combined (hybrid) and
rearrange themselves
to the shape and
orientation of a linear
shapeMolecular shape Linear
Angle between bond pair 1800
b) Whereas for hydrogen cyanide HCN both carbon atom and
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) y g y
nitrogen atom undergoes sp hybridisation in order to form a
linear structureExplanation
Energy level diagram for
carbon atom C
Energy level diagram for
nitrogen atom N
a) Ground state
Valence electronC 2s22p2
N 2s22p3
b) Excited state
One of the electron from
2s is promote to 2p
orbitals - equal energy
level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined
(hybrid) and form sp
hybridisation
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N
σσσσ
π
σσσσ
π
6 However there are a few limitation on valence bond theory such
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as when explaining the effect angle of bond-pair and bond-pair
electrons when there isare presence of lone pair electron in thecenter atom and also the difference of electronegativity
a) Bonding in ammonia NH3 and water H2O ~ Both NH3 and
H2O undergoes an arrangement similar to sp3 hybridisation
similar to that of C in methane Table below shows thehybridisation occur for nitrogen in ammonia and oxygen in water
Energy level diagram Energy level diagram
for nitrogen atom N for oxygen atom O
a) Ground state
Valence electron
N 2s22p3
O 2s2
2p4
b) Excited state
One of the electron
from 2s is promote to
2p orbitals
c) Hybridised state
One orbital of 2s and
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One orbital of 2s and
three orbitals from 2p
combined (hybrid) andrearrange themselves to
the shape and
orientation similar to that
of tetrahedral shape
Similar to arrangement in
tetrahedral
Similar to arrangement in
tetrahedral
of ammonia NH3 andwater H2O
Shape and angleShape trigonal pyramidal
Angle 1070
Shape bent
Angle 10450
Number of bond pair amp
lone pair electrons
Bond pair electron 3
Lone pair electron 1
Bond pair electron 2
Lone pair electron 2
bull The angle of bond pair - bond pair electrons in ammonia and0 0
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water are 1070 and 10450 respectively which is lesser than in
methane molecule (10950
) This can be explained by the factof the presence of lone-pair electrons in both ammonia and
water Since the lone pair - lone pair electron repulsion is
stronger than lone pair - bond pair electron repulsion than
bond pair - bond pair electron repulsion it is expected that therepulsion between lone pair - bond pair electrons in ammonia
is stronger than bond pair - bond pair electrons repulsion
hence caused the angle to squeeze to a smaller angle As
for water since there is a presence of lone pair - lone pair
electrons repulsion it results the bond pair - bond pair
electron repulsion to be much smaller hence caused the
angle to squeeze to a smaller angle Though valence bondtheory does not actually explain the hybridisation especially
when it relates to the repulsion occur involving lone pair
electron another theories shall be applied to study such
effect All these shall be discussed further in VSEPR theory
b) Effect of electronegativity and bonding angle between bond pair
- bond pair electrons ~
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- bond pair electrons
i As discussed earlier in the bonding of water (H2
O) the
molecular shape and angle is described in the diagram below
Sul hur which is also an element from Grou 16 formed h dro en
sulphide H2S when sulphur react with hydrogen However unlike
water the bonding angles is much smaller compare to water This is
due to the difference of electronegativity and also the bond length
between O and S in molecule Since O is more electronegative
compare to S the O-H bond is pulled closer toward O Furthermorethe bond length of O-H is shorter compare to S-H As a result the
bonding pair-bonding pair electrons repulse greater with each other in
H-O-H and caused the angle become greater
ii Another example is between NH3 and PH3 The orbital
f
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diagram for both NH3 and PH3 are described below Both
nitrogen N and phosphorous P are from the same groupwhich is Group 15
Similar to the above case P in phosphine is less
electronegative than N in ammonia and bond length of N-H
is shorter than P-H As a result H is pull closer to N and
repulsion between bonding of H-N-H is greater compare toH-P-H and caused the angle between H-N-H is greater
compare to H-P-H
324 Valence Shell Electron Pair Repulsion (VSEPR) Theory
1 M l l t i th th di i l t f t
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1 Molecular geometry is the three-dimensional arrangement of atoms
in a molecule A moleculersquos geometry affects its physical andchemical properties such as melting point boiling point density
and the types of reactions it undergoes The basic concept of
VSEPR theory is based on the three general rules below
a) As far as electron-pair repulsion is concerned double bonds andtriple bonds can be treated like single bonds This approximation is
good for qualitative purposes However you should realize that in
ldquo rdquo
because there are two or three bonds between two atoms theelectron density occupies more space
b) If a molecule has two or more resonance structures we can apply
the VSEPR model to any one of them Formal charges are usually
not shown
c) The order of repulsion strength of lone pair and bond pair are
lone-pair amp lone-pair electrons repulsion are the strongest
followed by lone-pair amp bond-pair electrons repulsion while
bond-pair amp bond-pair electrons repulsion is the weakest
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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3 Properties of Ionic Compound
a) Melting point - Ionic compound has giant ionic crystal lattice
which are hold by strong electrostatic attraction forces by
repeating of oppositely charged ions
Therefore the melting point of ionic compounds are usuallyvery high as a lot of energies are required to overcome the
electrostatic attraction forces and melted to form free moving
ions Therefore ionic compounds have very high melting
point
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b) Conductivity of electricity - Most ionic compounds do not
conduct electricity (insulator) in the solid state but do
conduct it when melted or when dissolved in water (exceptsome super-ionic conductors such as AgI which have
remarkable conductivity in the solid state) Solid ionic salt
consists of immobilized ions When it melts or dissolveshowever the ions are free to move and carry an electric
current
-
are hard (does not dent) rigid (does not bend) and brittle(cracks without deforming) These properties are due to the
powerful attractive forces that hold the ions in specific
positions throughout the crystal Moving the ions out of
position requires overcoming these forces so the sample
resists denting and bending If enough pressure is applied
ions of like charge are brought next to each other and
repulsive forces crack the ionic solid suddenly
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32 Covalent bondbull Studies of covalent bond was widely developed ever since
Lewis suggested that a chemical bond exist in a hydrogen gas
occur by sharing en electron between two hydrogen atoms
bull Electron pair that connect the 2 hydrogen atoms is called
covalent bond a bond in which two electrons are shared by
wo a oms an e e ec ron pa r a on e ween e wohydrogen atoms is also called as bonding pair electrons
bull In a covalent bond each electron in a shared pair is attracted to
the nuclei of both atoms This attraction holds the two atoms in
H2 together and is responsible for the formation of covalentbonds in other molecules
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2 A Lewis structure is a representation of covalent bonding
in which shared electron pairs are shown either as lines
or as pairs of dots between two atoms and lone pairsare shown as pairs of dots on individual atoms Only
valence electrons are shown in a Lewis structure
a) Consider the fluorine molecule F2 The electronconfiguration of F is 1s22s22p5 The 1s electrons are in
inner shell which is nearest to the nucleus For this reason
the do not artici ate in bond formation Thus each F
atom has seven valence electrons (2s22p5) Thereforeeach fluorine atom needed one electron to achieve octet
configuration (ns2np6)
F 1 s22s22p5 F 1 s22s22p5 1s2 2s2 2p6 1s2 2s2 2p6
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bull Oxygen atom has electronic configuration of 1s22s22p4 To
achieve stable octet configuration (ns2np6) each oxygen
atom need 2 electrons Hence when 2 oxygen atomsinteract they shared two electrons in between each other as
described in diagram below
O 1s2
2s2
2p4
O 1s2
2s2
2p4
1s2
2s2
2p6
1s2
2s2
2p6
bull From the structure of oxygen molecule formed each oxygen
atom shared two electrons from each other to form a double
bond in order to achieve octet configuration among each
oxygen atom
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bull Nitrogen has electronic configuration of 1s22s22p3 and
required 3 electrons to achieve octet configuration (ns2np6)
In this case each nitrogen atom shared 3 electrons fromeach of its atom to form triple bond
N 1 s22s22p3 N 1 s22s22p3 1s2 2s2 2p6 1s2 2s2 2p6
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Water H2O
H 1 s1 O 2 s22p4
Carbon dioxide CO2
C 2 s22p2 O 2 s22p4
Ammonia NH3
H 1 s1 N 2 s22p3Ethene C2H4
H 1 s1 C 2 s22p2
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Hydrogen cyanide HCN
H 1 s1 N 2 s22p3 C 2 s22p2
Ethanoic acid CH3COOH
H 1 s1 O 2 s22p4 C 2 s22p2
Tetrachloromethane CCl4C 2 s22p2 Cl 3s23p5
Ethyne C2H2
H 1 s1 C 2 s22p2
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a) Note that in ethene hydrogen cyanide and ethanoic acid all
the valence electrons are used in bonding there are no lone
pairs on the carbon atoms In fact most of the stablemolecules containing carbon do not have lone pairs on the
carbon atoms
b) Multiple bonds (double bond or triple bond) are shorter than single covalent bonds Bond length is defined as the
distance between the nuclei of two covalently bonded atoms
nitrogen triple bonds are shorter than double bonds whichin turn are shorter than single bonds The shorter multiple
bonds are also more stable than single bonds
c) Covalent bond not only exist in neutral molecule but also insome molecular ions Table below shows a few example of
molecular ions which have covalent bonds in its molecule
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Carbonate ion CO32-
C 2 s22p2 O 2 s22p4
Cyanide ion CN-
N 2 s22p3 C 2 s22p2
u p a e on 4-
S 3 s23p4 O 2 s22p4
ra e on 3-
N 2 s22p3 O 2 s22p4
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321 Exception of Octet Rules
1 From the Lewis structure sketched for sulphate ion SO42- we can
see that the center metal atom (sulphur) has more than 8electrons For this the molecule is described as molecule that can
expanded octet These molecules that have more than 8
electrons are located at Period 3 and below as these center
atoms have empty d-orbital available to expand the number ofelectrons positioned in the center atom Those center atom from
Period 2 such as C N O and F can only allocate a maximum
Example Both phosphorous and nitrogen are elements from Group15 with the valence electron of ns2np3 They can both form NCl3and PCl3 respectively when react with limited amount of chlorine
however under excess chlorine only PCl 5 can be formed but
not NCl 5 Explain the statement bolded
Solution This is due to phosphorous which is from period
3 have empty d-orbital to expand the octet However
nitrogen which is from Period 2 do not have emptyorbital and can only allocate 8 electrons in its shell
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Phosphorous pentachloride PCl5P 3 s23p3 Cl 3s23p5
Sulphur hexafluoride SF6
S 3 s23p4 F 2 s22p5
Bromine pentachloride BrCl5Cl 3s23p5 Br 4s24p5
Xenon tetrafluoride XeF4
Xe 5s25p6 F 2 s22p5
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2 There are also some stable covalent compounds which have
less than 8 electrons in their center atom (incomplete octet)
The center atoms are usually metals with great number ofvalence electrons with small atomic radius such as
beryllium boron and aluminium
BeCl2Be 2s2 Cl 3s23p5 BF3B 2 s22p1 F 2 s22p5 AlCl3Al 3s23p1 F 3 s23p5
These compound possessed stability due to their short bondlength between center atom with surround atom Furthermore
they can form resonance structure between the center atom
and surrounding atoms
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3 There are also some molecules which contain an odd
number of electrons Among the most common compounds
are nitrogen monoxide (NO) and nitrogen dioxide (NO2)
Odd-electron molecules are sometimes called radicals
Nitrogen monoxide NO
N 2 s22p3 O 2 s22p4
Nitrogen dioxide NO2
N 2 s22p3 O 2 s22p4
any ra ca s are g y reac ve e reason s t at t ere
is a tendency for the unpaired electron to form a
covalent bond with an unpaired electron on another
molecule For example when two nitrogen dioxide
molecules collide they form dinitrogen tetroxide
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322 Dative bond
1 As shown in the Lewis structure of nitrate ion one of the N-Obond is drawn as rarr The bond rarr placed is called as dative
bond (also known as coordinative bond) where dative
bond is defined as a covalent bond in which one of the atoms
donate the lone pair electrons available Although theproperties of a coordinate covalent bond do not differ from
those of a normal covalent bond because all electrons are
a e no ma er w a e r source
2 Dative bond is usually applied for these few circumstances
below
a) To assist atom molecule ion that not yet achieved octet
configuration Making use of atom which has lone pair
electrons to those which are lack of electrons
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Sulphur dioxide SO2Sulphur trioxide SO3 Carbon monoxide CO
Ozone molecule O3 Water with hydrogen ion
S
OO
O
Ammonia with hydrogen ion Ammonia with boron trifluoride
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b) Formation of dimer - In order for some compounds which
have incomplete octet to achieve stability they tend to form
dimer or polymer among themselves by using dative bondTwo of the most common examples are aluminium trichloride
and beryllium dichloride
Monomer of AlCl3Dimer of aluminium chloride Al
2
Cl6
Monomer of BeCl2 Dimer of Be2Cl4 Polymer of (BeCl2)n
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c) Formation of coordination compounds - Coordination
compounds are substances that contain at least one complex
ion a species consisting of a central metal cation (either atransition metal or a main-group metal) that is bonded to
molecules andor anions called ligands via dative
(coordinative) bond
hexaaquacopper (II) ion [Cu(H2O)6]2+ tetraamminenickel (II) ion [Ni(NH3)4]
2+
Hexacyanoferrate (III) ion [Fe(CN)6]3- Trioxalatocobaltate(III)ion [Co(C2O4)3]
3-
3 2 3 Hybridisation Theory
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323 Hybridisation Theory
Valence bond theory
bull The basic principle of valence bond theory is that a covalentbond is formed when orbitals of two atoms overlap and the
overlapped region which is between the nuclei is occupied by
a pair of electrons The central themes of valence bond theory
derive from this principle
ndash Opposing spins of the electron pair ~ Stated in Paulis
orbitals has a maximum capacity of two electrons that musthave opposite spins For example when a covalent bond is
formed in molecule of hydrogen H2 the two 1s electrons of
two H atoms occupy the overlapping 1s orbitals and have
opposite spins
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bull Maximum overlap of bonding orbitals ~ The bond strength
depends on attraction of the nuclei for the shared electrons
so the greater the orbital overlap the stronger (morestable) the bond The extent of overlap depends on the
shapes and directions of the orbitals An s orbital is spherical
but p and d orbitals have more electron density in one
direction than in another Thus whenever possible a bond
involving p or d orbitals will be oriented in the direction that
maximizes overlap For example in hydrogen fluoride (HF)
bond the 1s orbital of H overlaps the half-filled 2p orbital of Falong the long axis of that orbital Any other direction would
result in less overlap and thus a weaker bond
H b idi ti f t i bit l T t f th
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bull Hybridisation of atomic orbitals ~ To account for the
bonding in simple diatomic molecules like HF we picture the
direct overlap of s and p orbitals of isolated atoms But howcan we account for the shapes of so many molecules and
polyatomic ions through the overlap of spherical s orbitals
dumbbell-shaped p orbitals and cloverleaf-shaped d orbitals
Linus Pauling proposed that the valence atomic orbitals in the
molecule are different from those in the isolated atoms The
spatial orientations of these new orbitals lead to more stable
bonds and are consistent with observed molecular shapesThe process of orbital mixing is called hybridisation and
the new atomic orbitals are called hybrid orbitals Two key
points about the number and type of hybrid orbitals are that
i The number of hybrid orbitals obtained equals the number ofatomic orbitals mixed
ii The type of hybrid orbitals obtained varies with the types of
atomic orbitals mixed
3 2 3 1 T f h b idi ti
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3231 Type of hybridisation
1 sp3 hybridisation ~ When four electron groups surround
the central atom the center atom involved must prepare fourorbitals with equal energies to overlap with the four
surrounding electron groups Valence Bond theory uses
hypothetical hybrid orbitals which are atomic orbitals
obtained when two or more non-equivalent orbitals of the
same atom combine in preparation for covalent bond
formation H bridisation is the term a lied to the mixin of
atomic orbitals in an atom (usually a central atom) togenerate a set of hybrid orbitals In the case of sp3 We can
generate four equivalent hybrid orbitals from the center
atom by mixing the s orbital and the three p orbitals
Explanation Energy level diagram Diagram of orbitals
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Explanation Energy level diagram Diagram of orbitals
a) Ground state
Carbon which act as thecenter atom has the valence
electron of 2s22p2
b) Excited state One of the electron from 2s
is promote to 2p orbitals -
equa energy eve
c) Hybridised state
One orbital of 2s and three
orbitals of 2p combined(hybrid) and rearrange
themselves to the shape
and orientation of a
tetrahedral shape
Molecular shape Tetrahedral
Angle between bond pair 10950
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2 sp2 hybridisation ~ When three electron groups surround
the central atom the center atom involved must prepare three
orbitals with equal energies to overlap with the three
surrounding electron groups In sp2 hybridisation three
equivalent (in terms of energy level) from the center atom by
mixing the one s orbital and the two p orbitals Using boron
trifluoride (BF3) as example sp2
hybridisation is explained
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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3 sp hybridisation ~ When two electron groups surround
the central atom we observe a linear shape which
means that the bonding orbitals must have a linear
orientation VB theory explains this by proposing thatmixing two nonequivalent orbitals of a central atom one
s and one p gives rise to two equivalent sp hybrid
orbitals that lie 1800 apart
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Beryllium which act as the
center atom has the valence
electron of 2s2
b) Excited state
One of the electron from 2s ispromote to 2p orbitals - equal
energy level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of a
linear shape
Molecular shape Linear
Angle between bond pair 1800
4 The concept of hybridisation is also useful to explain molecules with
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p y p
doubletriple bonds By using the concept of the direct overlapping
orbitals and side-touch lapping orbitals the formation of multiple
bonds in ethene C2H4 and ethyne C2H2 are described
a) Ethene C2H4 - Hybridisation take place for both carbon atoms in
ethene molecule is sp2 hybridisation
bull When both hybridised C is bonded together one of the hybridisedorbital overlapped directly between each other while the other two
hybridised orbitals overlapped directly with hydrogen atoms
z
atoms form a side-touch bond between each other and form anotherbond as shown in the diagram below
bull From the diagram with C=C there are two types of bond A sigma-
bond (σ-bond) is covalent bonds formed by orbitals overlapping end-
to-end with the electron density concentrated between the nuclei ofthe bonding atoms while the second type is called a pi bond (π-
bond) which is defined as a covalent bond formed by sideways
overlapping orbitals with electron density concentrated above and
below the plane of the nuclei of the bonding atoms
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital of 2s and two
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of atrigonal planar shape
Note that on unhybridised pz
orbital an electron is presence
Molecular shape Trigonal planar Angle between bond pair 1200
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b) Ethyne C H has the Lewis structure of H CequivC H which the
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b) Ethyne C2H2 has the Lewis structure of HminusCequivCminusH which the
bonding can be explain using sp hybridisation Table below
described how the hybridisation take place on each carbonatom and how the formation of triple bond occur
c) From the diagram formation of -CequivC- is due to the formation
of one sigma-bond by direct overlapping of one of the twohybrid orbitals on each C while the other two pi-bond bonds
are formed as a result of side-lapping of the each two
y z
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital from 2s and 2p
orbitals combined (hybrid) and
rearrange themselves to the
shape and orientation of a
linear shapeNote that on unhybridised py amp
pz orbitals electrons are
presence to form two
π-bonds Molecular shape Linear Angle between bond pair 1800
sp
pz py
5 Other examples of applications in valence bond theory includes the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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formation of nitrogen molecule N2 and hydrogen cyanide HCN
molecule
a) Nitrogen gas is Earths most abundant gas as it cover 78 of the
content of our air Nitrogen molecule is an inert gas thanks to its
short covalent bond and also its strong triple bond Therefore
a lot of heats are required to break the chemical bond of nitrogenbefore it can be applied in industries Using valence bond theory
the bonding of nitrogen molecule is explained in the diagram
b) So when the two hybridised nitrogen atom interacting amongeach other they hence formed a linear shape and two pi-bonds
(πminusbonds) are formed as a result of side-lapping of unhybridised
py and pz orbital respectively
N N
ExplanationEnergy level
diagramDiagram of orbitals
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diagram
a) Ground state
Both nitrogen whichact as the center atom
has the valence
electron of 2s22p3
b) Excited state One of the electron
from 2s is promote to 2p
orbitals - e ual ener
level
c) Hybridised state
One orbital of 2s and
one orbitals of 2p
combined (hybrid) and
rearrange themselves
to the shape and
orientation of a linear
shapeMolecular shape Linear
Angle between bond pair 1800
b) Whereas for hydrogen cyanide HCN both carbon atom and
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) y g y
nitrogen atom undergoes sp hybridisation in order to form a
linear structureExplanation
Energy level diagram for
carbon atom C
Energy level diagram for
nitrogen atom N
a) Ground state
Valence electronC 2s22p2
N 2s22p3
b) Excited state
One of the electron from
2s is promote to 2p
orbitals - equal energy
level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined
(hybrid) and form sp
hybridisation
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N
σσσσ
π
σσσσ
π
6 However there are a few limitation on valence bond theory such
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as when explaining the effect angle of bond-pair and bond-pair
electrons when there isare presence of lone pair electron in thecenter atom and also the difference of electronegativity
a) Bonding in ammonia NH3 and water H2O ~ Both NH3 and
H2O undergoes an arrangement similar to sp3 hybridisation
similar to that of C in methane Table below shows thehybridisation occur for nitrogen in ammonia and oxygen in water
Energy level diagram Energy level diagram
for nitrogen atom N for oxygen atom O
a) Ground state
Valence electron
N 2s22p3
O 2s2
2p4
b) Excited state
One of the electron
from 2s is promote to
2p orbitals
c) Hybridised state
One orbital of 2s and
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One orbital of 2s and
three orbitals from 2p
combined (hybrid) andrearrange themselves to
the shape and
orientation similar to that
of tetrahedral shape
Similar to arrangement in
tetrahedral
Similar to arrangement in
tetrahedral
of ammonia NH3 andwater H2O
Shape and angleShape trigonal pyramidal
Angle 1070
Shape bent
Angle 10450
Number of bond pair amp
lone pair electrons
Bond pair electron 3
Lone pair electron 1
Bond pair electron 2
Lone pair electron 2
bull The angle of bond pair - bond pair electrons in ammonia and0 0
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water are 1070 and 10450 respectively which is lesser than in
methane molecule (10950
) This can be explained by the factof the presence of lone-pair electrons in both ammonia and
water Since the lone pair - lone pair electron repulsion is
stronger than lone pair - bond pair electron repulsion than
bond pair - bond pair electron repulsion it is expected that therepulsion between lone pair - bond pair electrons in ammonia
is stronger than bond pair - bond pair electrons repulsion
hence caused the angle to squeeze to a smaller angle As
for water since there is a presence of lone pair - lone pair
electrons repulsion it results the bond pair - bond pair
electron repulsion to be much smaller hence caused the
angle to squeeze to a smaller angle Though valence bondtheory does not actually explain the hybridisation especially
when it relates to the repulsion occur involving lone pair
electron another theories shall be applied to study such
effect All these shall be discussed further in VSEPR theory
b) Effect of electronegativity and bonding angle between bond pair
- bond pair electrons ~
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- bond pair electrons
i As discussed earlier in the bonding of water (H2
O) the
molecular shape and angle is described in the diagram below
Sul hur which is also an element from Grou 16 formed h dro en
sulphide H2S when sulphur react with hydrogen However unlike
water the bonding angles is much smaller compare to water This is
due to the difference of electronegativity and also the bond length
between O and S in molecule Since O is more electronegative
compare to S the O-H bond is pulled closer toward O Furthermorethe bond length of O-H is shorter compare to S-H As a result the
bonding pair-bonding pair electrons repulse greater with each other in
H-O-H and caused the angle become greater
ii Another example is between NH3 and PH3 The orbital
f
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diagram for both NH3 and PH3 are described below Both
nitrogen N and phosphorous P are from the same groupwhich is Group 15
Similar to the above case P in phosphine is less
electronegative than N in ammonia and bond length of N-H
is shorter than P-H As a result H is pull closer to N and
repulsion between bonding of H-N-H is greater compare toH-P-H and caused the angle between H-N-H is greater
compare to H-P-H
324 Valence Shell Electron Pair Repulsion (VSEPR) Theory
1 M l l t i th th di i l t f t
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1 Molecular geometry is the three-dimensional arrangement of atoms
in a molecule A moleculersquos geometry affects its physical andchemical properties such as melting point boiling point density
and the types of reactions it undergoes The basic concept of
VSEPR theory is based on the three general rules below
a) As far as electron-pair repulsion is concerned double bonds andtriple bonds can be treated like single bonds This approximation is
good for qualitative purposes However you should realize that in
ldquo rdquo
because there are two or three bonds between two atoms theelectron density occupies more space
b) If a molecule has two or more resonance structures we can apply
the VSEPR model to any one of them Formal charges are usually
not shown
c) The order of repulsion strength of lone pair and bond pair are
lone-pair amp lone-pair electrons repulsion are the strongest
followed by lone-pair amp bond-pair electrons repulsion while
bond-pair amp bond-pair electrons repulsion is the weakest
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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3 Properties of Ionic Compound
a) Melting point - Ionic compound has giant ionic crystal lattice
which are hold by strong electrostatic attraction forces by
repeating of oppositely charged ions
Therefore the melting point of ionic compounds are usuallyvery high as a lot of energies are required to overcome the
electrostatic attraction forces and melted to form free moving
ions Therefore ionic compounds have very high melting
point
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b) Conductivity of electricity - Most ionic compounds do not
conduct electricity (insulator) in the solid state but do
conduct it when melted or when dissolved in water (exceptsome super-ionic conductors such as AgI which have
remarkable conductivity in the solid state) Solid ionic salt
consists of immobilized ions When it melts or dissolveshowever the ions are free to move and carry an electric
current
-
are hard (does not dent) rigid (does not bend) and brittle(cracks without deforming) These properties are due to the
powerful attractive forces that hold the ions in specific
positions throughout the crystal Moving the ions out of
position requires overcoming these forces so the sample
resists denting and bending If enough pressure is applied
ions of like charge are brought next to each other and
repulsive forces crack the ionic solid suddenly
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32 Covalent bondbull Studies of covalent bond was widely developed ever since
Lewis suggested that a chemical bond exist in a hydrogen gas
occur by sharing en electron between two hydrogen atoms
bull Electron pair that connect the 2 hydrogen atoms is called
covalent bond a bond in which two electrons are shared by
wo a oms an e e ec ron pa r a on e ween e wohydrogen atoms is also called as bonding pair electrons
bull In a covalent bond each electron in a shared pair is attracted to
the nuclei of both atoms This attraction holds the two atoms in
H2 together and is responsible for the formation of covalentbonds in other molecules
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2 A Lewis structure is a representation of covalent bonding
in which shared electron pairs are shown either as lines
or as pairs of dots between two atoms and lone pairsare shown as pairs of dots on individual atoms Only
valence electrons are shown in a Lewis structure
a) Consider the fluorine molecule F2 The electronconfiguration of F is 1s22s22p5 The 1s electrons are in
inner shell which is nearest to the nucleus For this reason
the do not artici ate in bond formation Thus each F
atom has seven valence electrons (2s22p5) Thereforeeach fluorine atom needed one electron to achieve octet
configuration (ns2np6)
F 1 s22s22p5 F 1 s22s22p5 1s2 2s2 2p6 1s2 2s2 2p6
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bull Oxygen atom has electronic configuration of 1s22s22p4 To
achieve stable octet configuration (ns2np6) each oxygen
atom need 2 electrons Hence when 2 oxygen atomsinteract they shared two electrons in between each other as
described in diagram below
O 1s2
2s2
2p4
O 1s2
2s2
2p4
1s2
2s2
2p6
1s2
2s2
2p6
bull From the structure of oxygen molecule formed each oxygen
atom shared two electrons from each other to form a double
bond in order to achieve octet configuration among each
oxygen atom
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bull Nitrogen has electronic configuration of 1s22s22p3 and
required 3 electrons to achieve octet configuration (ns2np6)
In this case each nitrogen atom shared 3 electrons fromeach of its atom to form triple bond
N 1 s22s22p3 N 1 s22s22p3 1s2 2s2 2p6 1s2 2s2 2p6
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Water H2O
H 1 s1 O 2 s22p4
Carbon dioxide CO2
C 2 s22p2 O 2 s22p4
Ammonia NH3
H 1 s1 N 2 s22p3Ethene C2H4
H 1 s1 C 2 s22p2
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Hydrogen cyanide HCN
H 1 s1 N 2 s22p3 C 2 s22p2
Ethanoic acid CH3COOH
H 1 s1 O 2 s22p4 C 2 s22p2
Tetrachloromethane CCl4C 2 s22p2 Cl 3s23p5
Ethyne C2H2
H 1 s1 C 2 s22p2
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a) Note that in ethene hydrogen cyanide and ethanoic acid all
the valence electrons are used in bonding there are no lone
pairs on the carbon atoms In fact most of the stablemolecules containing carbon do not have lone pairs on the
carbon atoms
b) Multiple bonds (double bond or triple bond) are shorter than single covalent bonds Bond length is defined as the
distance between the nuclei of two covalently bonded atoms
nitrogen triple bonds are shorter than double bonds whichin turn are shorter than single bonds The shorter multiple
bonds are also more stable than single bonds
c) Covalent bond not only exist in neutral molecule but also insome molecular ions Table below shows a few example of
molecular ions which have covalent bonds in its molecule
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Carbonate ion CO32-
C 2 s22p2 O 2 s22p4
Cyanide ion CN-
N 2 s22p3 C 2 s22p2
u p a e on 4-
S 3 s23p4 O 2 s22p4
ra e on 3-
N 2 s22p3 O 2 s22p4
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321 Exception of Octet Rules
1 From the Lewis structure sketched for sulphate ion SO42- we can
see that the center metal atom (sulphur) has more than 8electrons For this the molecule is described as molecule that can
expanded octet These molecules that have more than 8
electrons are located at Period 3 and below as these center
atoms have empty d-orbital available to expand the number ofelectrons positioned in the center atom Those center atom from
Period 2 such as C N O and F can only allocate a maximum
Example Both phosphorous and nitrogen are elements from Group15 with the valence electron of ns2np3 They can both form NCl3and PCl3 respectively when react with limited amount of chlorine
however under excess chlorine only PCl 5 can be formed but
not NCl 5 Explain the statement bolded
Solution This is due to phosphorous which is from period
3 have empty d-orbital to expand the octet However
nitrogen which is from Period 2 do not have emptyorbital and can only allocate 8 electrons in its shell
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Phosphorous pentachloride PCl5P 3 s23p3 Cl 3s23p5
Sulphur hexafluoride SF6
S 3 s23p4 F 2 s22p5
Bromine pentachloride BrCl5Cl 3s23p5 Br 4s24p5
Xenon tetrafluoride XeF4
Xe 5s25p6 F 2 s22p5
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2 There are also some stable covalent compounds which have
less than 8 electrons in their center atom (incomplete octet)
The center atoms are usually metals with great number ofvalence electrons with small atomic radius such as
beryllium boron and aluminium
BeCl2Be 2s2 Cl 3s23p5 BF3B 2 s22p1 F 2 s22p5 AlCl3Al 3s23p1 F 3 s23p5
These compound possessed stability due to their short bondlength between center atom with surround atom Furthermore
they can form resonance structure between the center atom
and surrounding atoms
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3 There are also some molecules which contain an odd
number of electrons Among the most common compounds
are nitrogen monoxide (NO) and nitrogen dioxide (NO2)
Odd-electron molecules are sometimes called radicals
Nitrogen monoxide NO
N 2 s22p3 O 2 s22p4
Nitrogen dioxide NO2
N 2 s22p3 O 2 s22p4
any ra ca s are g y reac ve e reason s t at t ere
is a tendency for the unpaired electron to form a
covalent bond with an unpaired electron on another
molecule For example when two nitrogen dioxide
molecules collide they form dinitrogen tetroxide
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322 Dative bond
1 As shown in the Lewis structure of nitrate ion one of the N-Obond is drawn as rarr The bond rarr placed is called as dative
bond (also known as coordinative bond) where dative
bond is defined as a covalent bond in which one of the atoms
donate the lone pair electrons available Although theproperties of a coordinate covalent bond do not differ from
those of a normal covalent bond because all electrons are
a e no ma er w a e r source
2 Dative bond is usually applied for these few circumstances
below
a) To assist atom molecule ion that not yet achieved octet
configuration Making use of atom which has lone pair
electrons to those which are lack of electrons
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Sulphur dioxide SO2Sulphur trioxide SO3 Carbon monoxide CO
Ozone molecule O3 Water with hydrogen ion
S
OO
O
Ammonia with hydrogen ion Ammonia with boron trifluoride
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b) Formation of dimer - In order for some compounds which
have incomplete octet to achieve stability they tend to form
dimer or polymer among themselves by using dative bondTwo of the most common examples are aluminium trichloride
and beryllium dichloride
Monomer of AlCl3Dimer of aluminium chloride Al
2
Cl6
Monomer of BeCl2 Dimer of Be2Cl4 Polymer of (BeCl2)n
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c) Formation of coordination compounds - Coordination
compounds are substances that contain at least one complex
ion a species consisting of a central metal cation (either atransition metal or a main-group metal) that is bonded to
molecules andor anions called ligands via dative
(coordinative) bond
hexaaquacopper (II) ion [Cu(H2O)6]2+ tetraamminenickel (II) ion [Ni(NH3)4]
2+
Hexacyanoferrate (III) ion [Fe(CN)6]3- Trioxalatocobaltate(III)ion [Co(C2O4)3]
3-
3 2 3 Hybridisation Theory
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323 Hybridisation Theory
Valence bond theory
bull The basic principle of valence bond theory is that a covalentbond is formed when orbitals of two atoms overlap and the
overlapped region which is between the nuclei is occupied by
a pair of electrons The central themes of valence bond theory
derive from this principle
ndash Opposing spins of the electron pair ~ Stated in Paulis
orbitals has a maximum capacity of two electrons that musthave opposite spins For example when a covalent bond is
formed in molecule of hydrogen H2 the two 1s electrons of
two H atoms occupy the overlapping 1s orbitals and have
opposite spins
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bull Maximum overlap of bonding orbitals ~ The bond strength
depends on attraction of the nuclei for the shared electrons
so the greater the orbital overlap the stronger (morestable) the bond The extent of overlap depends on the
shapes and directions of the orbitals An s orbital is spherical
but p and d orbitals have more electron density in one
direction than in another Thus whenever possible a bond
involving p or d orbitals will be oriented in the direction that
maximizes overlap For example in hydrogen fluoride (HF)
bond the 1s orbital of H overlaps the half-filled 2p orbital of Falong the long axis of that orbital Any other direction would
result in less overlap and thus a weaker bond
H b idi ti f t i bit l T t f th
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bull Hybridisation of atomic orbitals ~ To account for the
bonding in simple diatomic molecules like HF we picture the
direct overlap of s and p orbitals of isolated atoms But howcan we account for the shapes of so many molecules and
polyatomic ions through the overlap of spherical s orbitals
dumbbell-shaped p orbitals and cloverleaf-shaped d orbitals
Linus Pauling proposed that the valence atomic orbitals in the
molecule are different from those in the isolated atoms The
spatial orientations of these new orbitals lead to more stable
bonds and are consistent with observed molecular shapesThe process of orbital mixing is called hybridisation and
the new atomic orbitals are called hybrid orbitals Two key
points about the number and type of hybrid orbitals are that
i The number of hybrid orbitals obtained equals the number ofatomic orbitals mixed
ii The type of hybrid orbitals obtained varies with the types of
atomic orbitals mixed
3 2 3 1 T f h b idi ti
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3231 Type of hybridisation
1 sp3 hybridisation ~ When four electron groups surround
the central atom the center atom involved must prepare fourorbitals with equal energies to overlap with the four
surrounding electron groups Valence Bond theory uses
hypothetical hybrid orbitals which are atomic orbitals
obtained when two or more non-equivalent orbitals of the
same atom combine in preparation for covalent bond
formation H bridisation is the term a lied to the mixin of
atomic orbitals in an atom (usually a central atom) togenerate a set of hybrid orbitals In the case of sp3 We can
generate four equivalent hybrid orbitals from the center
atom by mixing the s orbital and the three p orbitals
Explanation Energy level diagram Diagram of orbitals
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Explanation Energy level diagram Diagram of orbitals
a) Ground state
Carbon which act as thecenter atom has the valence
electron of 2s22p2
b) Excited state One of the electron from 2s
is promote to 2p orbitals -
equa energy eve
c) Hybridised state
One orbital of 2s and three
orbitals of 2p combined(hybrid) and rearrange
themselves to the shape
and orientation of a
tetrahedral shape
Molecular shape Tetrahedral
Angle between bond pair 10950
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2 sp2 hybridisation ~ When three electron groups surround
the central atom the center atom involved must prepare three
orbitals with equal energies to overlap with the three
surrounding electron groups In sp2 hybridisation three
equivalent (in terms of energy level) from the center atom by
mixing the one s orbital and the two p orbitals Using boron
trifluoride (BF3) as example sp2
hybridisation is explained
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3 sp hybridisation ~ When two electron groups surround
the central atom we observe a linear shape which
means that the bonding orbitals must have a linear
orientation VB theory explains this by proposing thatmixing two nonequivalent orbitals of a central atom one
s and one p gives rise to two equivalent sp hybrid
orbitals that lie 1800 apart
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Beryllium which act as the
center atom has the valence
electron of 2s2
b) Excited state
One of the electron from 2s ispromote to 2p orbitals - equal
energy level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of a
linear shape
Molecular shape Linear
Angle between bond pair 1800
4 The concept of hybridisation is also useful to explain molecules with
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p y p
doubletriple bonds By using the concept of the direct overlapping
orbitals and side-touch lapping orbitals the formation of multiple
bonds in ethene C2H4 and ethyne C2H2 are described
a) Ethene C2H4 - Hybridisation take place for both carbon atoms in
ethene molecule is sp2 hybridisation
bull When both hybridised C is bonded together one of the hybridisedorbital overlapped directly between each other while the other two
hybridised orbitals overlapped directly with hydrogen atoms
z
atoms form a side-touch bond between each other and form anotherbond as shown in the diagram below
bull From the diagram with C=C there are two types of bond A sigma-
bond (σ-bond) is covalent bonds formed by orbitals overlapping end-
to-end with the electron density concentrated between the nuclei ofthe bonding atoms while the second type is called a pi bond (π-
bond) which is defined as a covalent bond formed by sideways
overlapping orbitals with electron density concentrated above and
below the plane of the nuclei of the bonding atoms
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital of 2s and two
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of atrigonal planar shape
Note that on unhybridised pz
orbital an electron is presence
Molecular shape Trigonal planar Angle between bond pair 1200
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b) Ethyne C H has the Lewis structure of H CequivC H which the
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b) Ethyne C2H2 has the Lewis structure of HminusCequivCminusH which the
bonding can be explain using sp hybridisation Table below
described how the hybridisation take place on each carbonatom and how the formation of triple bond occur
c) From the diagram formation of -CequivC- is due to the formation
of one sigma-bond by direct overlapping of one of the twohybrid orbitals on each C while the other two pi-bond bonds
are formed as a result of side-lapping of the each two
y z
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital from 2s and 2p
orbitals combined (hybrid) and
rearrange themselves to the
shape and orientation of a
linear shapeNote that on unhybridised py amp
pz orbitals electrons are
presence to form two
π-bonds Molecular shape Linear Angle between bond pair 1800
sp
pz py
5 Other examples of applications in valence bond theory includes the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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formation of nitrogen molecule N2 and hydrogen cyanide HCN
molecule
a) Nitrogen gas is Earths most abundant gas as it cover 78 of the
content of our air Nitrogen molecule is an inert gas thanks to its
short covalent bond and also its strong triple bond Therefore
a lot of heats are required to break the chemical bond of nitrogenbefore it can be applied in industries Using valence bond theory
the bonding of nitrogen molecule is explained in the diagram
b) So when the two hybridised nitrogen atom interacting amongeach other they hence formed a linear shape and two pi-bonds
(πminusbonds) are formed as a result of side-lapping of unhybridised
py and pz orbital respectively
N N
ExplanationEnergy level
diagramDiagram of orbitals
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diagram
a) Ground state
Both nitrogen whichact as the center atom
has the valence
electron of 2s22p3
b) Excited state One of the electron
from 2s is promote to 2p
orbitals - e ual ener
level
c) Hybridised state
One orbital of 2s and
one orbitals of 2p
combined (hybrid) and
rearrange themselves
to the shape and
orientation of a linear
shapeMolecular shape Linear
Angle between bond pair 1800
b) Whereas for hydrogen cyanide HCN both carbon atom and
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) y g y
nitrogen atom undergoes sp hybridisation in order to form a
linear structureExplanation
Energy level diagram for
carbon atom C
Energy level diagram for
nitrogen atom N
a) Ground state
Valence electronC 2s22p2
N 2s22p3
b) Excited state
One of the electron from
2s is promote to 2p
orbitals - equal energy
level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined
(hybrid) and form sp
hybridisation
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N
σσσσ
π
σσσσ
π
6 However there are a few limitation on valence bond theory such
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as when explaining the effect angle of bond-pair and bond-pair
electrons when there isare presence of lone pair electron in thecenter atom and also the difference of electronegativity
a) Bonding in ammonia NH3 and water H2O ~ Both NH3 and
H2O undergoes an arrangement similar to sp3 hybridisation
similar to that of C in methane Table below shows thehybridisation occur for nitrogen in ammonia and oxygen in water
Energy level diagram Energy level diagram
for nitrogen atom N for oxygen atom O
a) Ground state
Valence electron
N 2s22p3
O 2s2
2p4
b) Excited state
One of the electron
from 2s is promote to
2p orbitals
c) Hybridised state
One orbital of 2s and
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One orbital of 2s and
three orbitals from 2p
combined (hybrid) andrearrange themselves to
the shape and
orientation similar to that
of tetrahedral shape
Similar to arrangement in
tetrahedral
Similar to arrangement in
tetrahedral
of ammonia NH3 andwater H2O
Shape and angleShape trigonal pyramidal
Angle 1070
Shape bent
Angle 10450
Number of bond pair amp
lone pair electrons
Bond pair electron 3
Lone pair electron 1
Bond pair electron 2
Lone pair electron 2
bull The angle of bond pair - bond pair electrons in ammonia and0 0
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water are 1070 and 10450 respectively which is lesser than in
methane molecule (10950
) This can be explained by the factof the presence of lone-pair electrons in both ammonia and
water Since the lone pair - lone pair electron repulsion is
stronger than lone pair - bond pair electron repulsion than
bond pair - bond pair electron repulsion it is expected that therepulsion between lone pair - bond pair electrons in ammonia
is stronger than bond pair - bond pair electrons repulsion
hence caused the angle to squeeze to a smaller angle As
for water since there is a presence of lone pair - lone pair
electrons repulsion it results the bond pair - bond pair
electron repulsion to be much smaller hence caused the
angle to squeeze to a smaller angle Though valence bondtheory does not actually explain the hybridisation especially
when it relates to the repulsion occur involving lone pair
electron another theories shall be applied to study such
effect All these shall be discussed further in VSEPR theory
b) Effect of electronegativity and bonding angle between bond pair
- bond pair electrons ~
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- bond pair electrons
i As discussed earlier in the bonding of water (H2
O) the
molecular shape and angle is described in the diagram below
Sul hur which is also an element from Grou 16 formed h dro en
sulphide H2S when sulphur react with hydrogen However unlike
water the bonding angles is much smaller compare to water This is
due to the difference of electronegativity and also the bond length
between O and S in molecule Since O is more electronegative
compare to S the O-H bond is pulled closer toward O Furthermorethe bond length of O-H is shorter compare to S-H As a result the
bonding pair-bonding pair electrons repulse greater with each other in
H-O-H and caused the angle become greater
ii Another example is between NH3 and PH3 The orbital
f
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diagram for both NH3 and PH3 are described below Both
nitrogen N and phosphorous P are from the same groupwhich is Group 15
Similar to the above case P in phosphine is less
electronegative than N in ammonia and bond length of N-H
is shorter than P-H As a result H is pull closer to N and
repulsion between bonding of H-N-H is greater compare toH-P-H and caused the angle between H-N-H is greater
compare to H-P-H
324 Valence Shell Electron Pair Repulsion (VSEPR) Theory
1 M l l t i th th di i l t f t
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1 Molecular geometry is the three-dimensional arrangement of atoms
in a molecule A moleculersquos geometry affects its physical andchemical properties such as melting point boiling point density
and the types of reactions it undergoes The basic concept of
VSEPR theory is based on the three general rules below
a) As far as electron-pair repulsion is concerned double bonds andtriple bonds can be treated like single bonds This approximation is
good for qualitative purposes However you should realize that in
ldquo rdquo
because there are two or three bonds between two atoms theelectron density occupies more space
b) If a molecule has two or more resonance structures we can apply
the VSEPR model to any one of them Formal charges are usually
not shown
c) The order of repulsion strength of lone pair and bond pair are
lone-pair amp lone-pair electrons repulsion are the strongest
followed by lone-pair amp bond-pair electrons repulsion while
bond-pair amp bond-pair electrons repulsion is the weakest
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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b) Conductivity of electricity - Most ionic compounds do not
conduct electricity (insulator) in the solid state but do
conduct it when melted or when dissolved in water (exceptsome super-ionic conductors such as AgI which have
remarkable conductivity in the solid state) Solid ionic salt
consists of immobilized ions When it melts or dissolveshowever the ions are free to move and carry an electric
current
-
are hard (does not dent) rigid (does not bend) and brittle(cracks without deforming) These properties are due to the
powerful attractive forces that hold the ions in specific
positions throughout the crystal Moving the ions out of
position requires overcoming these forces so the sample
resists denting and bending If enough pressure is applied
ions of like charge are brought next to each other and
repulsive forces crack the ionic solid suddenly
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32 Covalent bondbull Studies of covalent bond was widely developed ever since
Lewis suggested that a chemical bond exist in a hydrogen gas
occur by sharing en electron between two hydrogen atoms
bull Electron pair that connect the 2 hydrogen atoms is called
covalent bond a bond in which two electrons are shared by
wo a oms an e e ec ron pa r a on e ween e wohydrogen atoms is also called as bonding pair electrons
bull In a covalent bond each electron in a shared pair is attracted to
the nuclei of both atoms This attraction holds the two atoms in
H2 together and is responsible for the formation of covalentbonds in other molecules
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2 A Lewis structure is a representation of covalent bonding
in which shared electron pairs are shown either as lines
or as pairs of dots between two atoms and lone pairsare shown as pairs of dots on individual atoms Only
valence electrons are shown in a Lewis structure
a) Consider the fluorine molecule F2 The electronconfiguration of F is 1s22s22p5 The 1s electrons are in
inner shell which is nearest to the nucleus For this reason
the do not artici ate in bond formation Thus each F
atom has seven valence electrons (2s22p5) Thereforeeach fluorine atom needed one electron to achieve octet
configuration (ns2np6)
F 1 s22s22p5 F 1 s22s22p5 1s2 2s2 2p6 1s2 2s2 2p6
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bull Oxygen atom has electronic configuration of 1s22s22p4 To
achieve stable octet configuration (ns2np6) each oxygen
atom need 2 electrons Hence when 2 oxygen atomsinteract they shared two electrons in between each other as
described in diagram below
O 1s2
2s2
2p4
O 1s2
2s2
2p4
1s2
2s2
2p6
1s2
2s2
2p6
bull From the structure of oxygen molecule formed each oxygen
atom shared two electrons from each other to form a double
bond in order to achieve octet configuration among each
oxygen atom
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bull Nitrogen has electronic configuration of 1s22s22p3 and
required 3 electrons to achieve octet configuration (ns2np6)
In this case each nitrogen atom shared 3 electrons fromeach of its atom to form triple bond
N 1 s22s22p3 N 1 s22s22p3 1s2 2s2 2p6 1s2 2s2 2p6
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Water H2O
H 1 s1 O 2 s22p4
Carbon dioxide CO2
C 2 s22p2 O 2 s22p4
Ammonia NH3
H 1 s1 N 2 s22p3Ethene C2H4
H 1 s1 C 2 s22p2
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Hydrogen cyanide HCN
H 1 s1 N 2 s22p3 C 2 s22p2
Ethanoic acid CH3COOH
H 1 s1 O 2 s22p4 C 2 s22p2
Tetrachloromethane CCl4C 2 s22p2 Cl 3s23p5
Ethyne C2H2
H 1 s1 C 2 s22p2
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a) Note that in ethene hydrogen cyanide and ethanoic acid all
the valence electrons are used in bonding there are no lone
pairs on the carbon atoms In fact most of the stablemolecules containing carbon do not have lone pairs on the
carbon atoms
b) Multiple bonds (double bond or triple bond) are shorter than single covalent bonds Bond length is defined as the
distance between the nuclei of two covalently bonded atoms
nitrogen triple bonds are shorter than double bonds whichin turn are shorter than single bonds The shorter multiple
bonds are also more stable than single bonds
c) Covalent bond not only exist in neutral molecule but also insome molecular ions Table below shows a few example of
molecular ions which have covalent bonds in its molecule
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Carbonate ion CO32-
C 2 s22p2 O 2 s22p4
Cyanide ion CN-
N 2 s22p3 C 2 s22p2
u p a e on 4-
S 3 s23p4 O 2 s22p4
ra e on 3-
N 2 s22p3 O 2 s22p4
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321 Exception of Octet Rules
1 From the Lewis structure sketched for sulphate ion SO42- we can
see that the center metal atom (sulphur) has more than 8electrons For this the molecule is described as molecule that can
expanded octet These molecules that have more than 8
electrons are located at Period 3 and below as these center
atoms have empty d-orbital available to expand the number ofelectrons positioned in the center atom Those center atom from
Period 2 such as C N O and F can only allocate a maximum
Example Both phosphorous and nitrogen are elements from Group15 with the valence electron of ns2np3 They can both form NCl3and PCl3 respectively when react with limited amount of chlorine
however under excess chlorine only PCl 5 can be formed but
not NCl 5 Explain the statement bolded
Solution This is due to phosphorous which is from period
3 have empty d-orbital to expand the octet However
nitrogen which is from Period 2 do not have emptyorbital and can only allocate 8 electrons in its shell
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Phosphorous pentachloride PCl5P 3 s23p3 Cl 3s23p5
Sulphur hexafluoride SF6
S 3 s23p4 F 2 s22p5
Bromine pentachloride BrCl5Cl 3s23p5 Br 4s24p5
Xenon tetrafluoride XeF4
Xe 5s25p6 F 2 s22p5
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2 There are also some stable covalent compounds which have
less than 8 electrons in their center atom (incomplete octet)
The center atoms are usually metals with great number ofvalence electrons with small atomic radius such as
beryllium boron and aluminium
BeCl2Be 2s2 Cl 3s23p5 BF3B 2 s22p1 F 2 s22p5 AlCl3Al 3s23p1 F 3 s23p5
These compound possessed stability due to their short bondlength between center atom with surround atom Furthermore
they can form resonance structure between the center atom
and surrounding atoms
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3 There are also some molecules which contain an odd
number of electrons Among the most common compounds
are nitrogen monoxide (NO) and nitrogen dioxide (NO2)
Odd-electron molecules are sometimes called radicals
Nitrogen monoxide NO
N 2 s22p3 O 2 s22p4
Nitrogen dioxide NO2
N 2 s22p3 O 2 s22p4
any ra ca s are g y reac ve e reason s t at t ere
is a tendency for the unpaired electron to form a
covalent bond with an unpaired electron on another
molecule For example when two nitrogen dioxide
molecules collide they form dinitrogen tetroxide
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322 Dative bond
1 As shown in the Lewis structure of nitrate ion one of the N-Obond is drawn as rarr The bond rarr placed is called as dative
bond (also known as coordinative bond) where dative
bond is defined as a covalent bond in which one of the atoms
donate the lone pair electrons available Although theproperties of a coordinate covalent bond do not differ from
those of a normal covalent bond because all electrons are
a e no ma er w a e r source
2 Dative bond is usually applied for these few circumstances
below
a) To assist atom molecule ion that not yet achieved octet
configuration Making use of atom which has lone pair
electrons to those which are lack of electrons
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Sulphur dioxide SO2Sulphur trioxide SO3 Carbon monoxide CO
Ozone molecule O3 Water with hydrogen ion
S
OO
O
Ammonia with hydrogen ion Ammonia with boron trifluoride
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b) Formation of dimer - In order for some compounds which
have incomplete octet to achieve stability they tend to form
dimer or polymer among themselves by using dative bondTwo of the most common examples are aluminium trichloride
and beryllium dichloride
Monomer of AlCl3Dimer of aluminium chloride Al
2
Cl6
Monomer of BeCl2 Dimer of Be2Cl4 Polymer of (BeCl2)n
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c) Formation of coordination compounds - Coordination
compounds are substances that contain at least one complex
ion a species consisting of a central metal cation (either atransition metal or a main-group metal) that is bonded to
molecules andor anions called ligands via dative
(coordinative) bond
hexaaquacopper (II) ion [Cu(H2O)6]2+ tetraamminenickel (II) ion [Ni(NH3)4]
2+
Hexacyanoferrate (III) ion [Fe(CN)6]3- Trioxalatocobaltate(III)ion [Co(C2O4)3]
3-
3 2 3 Hybridisation Theory
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323 Hybridisation Theory
Valence bond theory
bull The basic principle of valence bond theory is that a covalentbond is formed when orbitals of two atoms overlap and the
overlapped region which is between the nuclei is occupied by
a pair of electrons The central themes of valence bond theory
derive from this principle
ndash Opposing spins of the electron pair ~ Stated in Paulis
orbitals has a maximum capacity of two electrons that musthave opposite spins For example when a covalent bond is
formed in molecule of hydrogen H2 the two 1s electrons of
two H atoms occupy the overlapping 1s orbitals and have
opposite spins
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bull Maximum overlap of bonding orbitals ~ The bond strength
depends on attraction of the nuclei for the shared electrons
so the greater the orbital overlap the stronger (morestable) the bond The extent of overlap depends on the
shapes and directions of the orbitals An s orbital is spherical
but p and d orbitals have more electron density in one
direction than in another Thus whenever possible a bond
involving p or d orbitals will be oriented in the direction that
maximizes overlap For example in hydrogen fluoride (HF)
bond the 1s orbital of H overlaps the half-filled 2p orbital of Falong the long axis of that orbital Any other direction would
result in less overlap and thus a weaker bond
H b idi ti f t i bit l T t f th
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bull Hybridisation of atomic orbitals ~ To account for the
bonding in simple diatomic molecules like HF we picture the
direct overlap of s and p orbitals of isolated atoms But howcan we account for the shapes of so many molecules and
polyatomic ions through the overlap of spherical s orbitals
dumbbell-shaped p orbitals and cloverleaf-shaped d orbitals
Linus Pauling proposed that the valence atomic orbitals in the
molecule are different from those in the isolated atoms The
spatial orientations of these new orbitals lead to more stable
bonds and are consistent with observed molecular shapesThe process of orbital mixing is called hybridisation and
the new atomic orbitals are called hybrid orbitals Two key
points about the number and type of hybrid orbitals are that
i The number of hybrid orbitals obtained equals the number ofatomic orbitals mixed
ii The type of hybrid orbitals obtained varies with the types of
atomic orbitals mixed
3 2 3 1 T f h b idi ti
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3231 Type of hybridisation
1 sp3 hybridisation ~ When four electron groups surround
the central atom the center atom involved must prepare fourorbitals with equal energies to overlap with the four
surrounding electron groups Valence Bond theory uses
hypothetical hybrid orbitals which are atomic orbitals
obtained when two or more non-equivalent orbitals of the
same atom combine in preparation for covalent bond
formation H bridisation is the term a lied to the mixin of
atomic orbitals in an atom (usually a central atom) togenerate a set of hybrid orbitals In the case of sp3 We can
generate four equivalent hybrid orbitals from the center
atom by mixing the s orbital and the three p orbitals
Explanation Energy level diagram Diagram of orbitals
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Explanation Energy level diagram Diagram of orbitals
a) Ground state
Carbon which act as thecenter atom has the valence
electron of 2s22p2
b) Excited state One of the electron from 2s
is promote to 2p orbitals -
equa energy eve
c) Hybridised state
One orbital of 2s and three
orbitals of 2p combined(hybrid) and rearrange
themselves to the shape
and orientation of a
tetrahedral shape
Molecular shape Tetrahedral
Angle between bond pair 10950
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2 sp2 hybridisation ~ When three electron groups surround
the central atom the center atom involved must prepare three
orbitals with equal energies to overlap with the three
surrounding electron groups In sp2 hybridisation three
equivalent (in terms of energy level) from the center atom by
mixing the one s orbital and the two p orbitals Using boron
trifluoride (BF3) as example sp2
hybridisation is explained
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3 sp hybridisation ~ When two electron groups surround
the central atom we observe a linear shape which
means that the bonding orbitals must have a linear
orientation VB theory explains this by proposing thatmixing two nonequivalent orbitals of a central atom one
s and one p gives rise to two equivalent sp hybrid
orbitals that lie 1800 apart
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Beryllium which act as the
center atom has the valence
electron of 2s2
b) Excited state
One of the electron from 2s ispromote to 2p orbitals - equal
energy level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of a
linear shape
Molecular shape Linear
Angle between bond pair 1800
4 The concept of hybridisation is also useful to explain molecules with
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p y p
doubletriple bonds By using the concept of the direct overlapping
orbitals and side-touch lapping orbitals the formation of multiple
bonds in ethene C2H4 and ethyne C2H2 are described
a) Ethene C2H4 - Hybridisation take place for both carbon atoms in
ethene molecule is sp2 hybridisation
bull When both hybridised C is bonded together one of the hybridisedorbital overlapped directly between each other while the other two
hybridised orbitals overlapped directly with hydrogen atoms
z
atoms form a side-touch bond between each other and form anotherbond as shown in the diagram below
bull From the diagram with C=C there are two types of bond A sigma-
bond (σ-bond) is covalent bonds formed by orbitals overlapping end-
to-end with the electron density concentrated between the nuclei ofthe bonding atoms while the second type is called a pi bond (π-
bond) which is defined as a covalent bond formed by sideways
overlapping orbitals with electron density concentrated above and
below the plane of the nuclei of the bonding atoms
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital of 2s and two
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of atrigonal planar shape
Note that on unhybridised pz
orbital an electron is presence
Molecular shape Trigonal planar Angle between bond pair 1200
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b) Ethyne C H has the Lewis structure of H CequivC H which the
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b) Ethyne C2H2 has the Lewis structure of HminusCequivCminusH which the
bonding can be explain using sp hybridisation Table below
described how the hybridisation take place on each carbonatom and how the formation of triple bond occur
c) From the diagram formation of -CequivC- is due to the formation
of one sigma-bond by direct overlapping of one of the twohybrid orbitals on each C while the other two pi-bond bonds
are formed as a result of side-lapping of the each two
y z
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital from 2s and 2p
orbitals combined (hybrid) and
rearrange themselves to the
shape and orientation of a
linear shapeNote that on unhybridised py amp
pz orbitals electrons are
presence to form two
π-bonds Molecular shape Linear Angle between bond pair 1800
sp
pz py
5 Other examples of applications in valence bond theory includes the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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formation of nitrogen molecule N2 and hydrogen cyanide HCN
molecule
a) Nitrogen gas is Earths most abundant gas as it cover 78 of the
content of our air Nitrogen molecule is an inert gas thanks to its
short covalent bond and also its strong triple bond Therefore
a lot of heats are required to break the chemical bond of nitrogenbefore it can be applied in industries Using valence bond theory
the bonding of nitrogen molecule is explained in the diagram
b) So when the two hybridised nitrogen atom interacting amongeach other they hence formed a linear shape and two pi-bonds
(πminusbonds) are formed as a result of side-lapping of unhybridised
py and pz orbital respectively
N N
ExplanationEnergy level
diagramDiagram of orbitals
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diagram
a) Ground state
Both nitrogen whichact as the center atom
has the valence
electron of 2s22p3
b) Excited state One of the electron
from 2s is promote to 2p
orbitals - e ual ener
level
c) Hybridised state
One orbital of 2s and
one orbitals of 2p
combined (hybrid) and
rearrange themselves
to the shape and
orientation of a linear
shapeMolecular shape Linear
Angle between bond pair 1800
b) Whereas for hydrogen cyanide HCN both carbon atom and
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) y g y
nitrogen atom undergoes sp hybridisation in order to form a
linear structureExplanation
Energy level diagram for
carbon atom C
Energy level diagram for
nitrogen atom N
a) Ground state
Valence electronC 2s22p2
N 2s22p3
b) Excited state
One of the electron from
2s is promote to 2p
orbitals - equal energy
level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined
(hybrid) and form sp
hybridisation
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N
σσσσ
π
σσσσ
π
6 However there are a few limitation on valence bond theory such
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as when explaining the effect angle of bond-pair and bond-pair
electrons when there isare presence of lone pair electron in thecenter atom and also the difference of electronegativity
a) Bonding in ammonia NH3 and water H2O ~ Both NH3 and
H2O undergoes an arrangement similar to sp3 hybridisation
similar to that of C in methane Table below shows thehybridisation occur for nitrogen in ammonia and oxygen in water
Energy level diagram Energy level diagram
for nitrogen atom N for oxygen atom O
a) Ground state
Valence electron
N 2s22p3
O 2s2
2p4
b) Excited state
One of the electron
from 2s is promote to
2p orbitals
c) Hybridised state
One orbital of 2s and
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One orbital of 2s and
three orbitals from 2p
combined (hybrid) andrearrange themselves to
the shape and
orientation similar to that
of tetrahedral shape
Similar to arrangement in
tetrahedral
Similar to arrangement in
tetrahedral
of ammonia NH3 andwater H2O
Shape and angleShape trigonal pyramidal
Angle 1070
Shape bent
Angle 10450
Number of bond pair amp
lone pair electrons
Bond pair electron 3
Lone pair electron 1
Bond pair electron 2
Lone pair electron 2
bull The angle of bond pair - bond pair electrons in ammonia and0 0
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water are 1070 and 10450 respectively which is lesser than in
methane molecule (10950
) This can be explained by the factof the presence of lone-pair electrons in both ammonia and
water Since the lone pair - lone pair electron repulsion is
stronger than lone pair - bond pair electron repulsion than
bond pair - bond pair electron repulsion it is expected that therepulsion between lone pair - bond pair electrons in ammonia
is stronger than bond pair - bond pair electrons repulsion
hence caused the angle to squeeze to a smaller angle As
for water since there is a presence of lone pair - lone pair
electrons repulsion it results the bond pair - bond pair
electron repulsion to be much smaller hence caused the
angle to squeeze to a smaller angle Though valence bondtheory does not actually explain the hybridisation especially
when it relates to the repulsion occur involving lone pair
electron another theories shall be applied to study such
effect All these shall be discussed further in VSEPR theory
b) Effect of electronegativity and bonding angle between bond pair
- bond pair electrons ~
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- bond pair electrons
i As discussed earlier in the bonding of water (H2
O) the
molecular shape and angle is described in the diagram below
Sul hur which is also an element from Grou 16 formed h dro en
sulphide H2S when sulphur react with hydrogen However unlike
water the bonding angles is much smaller compare to water This is
due to the difference of electronegativity and also the bond length
between O and S in molecule Since O is more electronegative
compare to S the O-H bond is pulled closer toward O Furthermorethe bond length of O-H is shorter compare to S-H As a result the
bonding pair-bonding pair electrons repulse greater with each other in
H-O-H and caused the angle become greater
ii Another example is between NH3 and PH3 The orbital
f
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diagram for both NH3 and PH3 are described below Both
nitrogen N and phosphorous P are from the same groupwhich is Group 15
Similar to the above case P in phosphine is less
electronegative than N in ammonia and bond length of N-H
is shorter than P-H As a result H is pull closer to N and
repulsion between bonding of H-N-H is greater compare toH-P-H and caused the angle between H-N-H is greater
compare to H-P-H
324 Valence Shell Electron Pair Repulsion (VSEPR) Theory
1 M l l t i th th di i l t f t
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1 Molecular geometry is the three-dimensional arrangement of atoms
in a molecule A moleculersquos geometry affects its physical andchemical properties such as melting point boiling point density
and the types of reactions it undergoes The basic concept of
VSEPR theory is based on the three general rules below
a) As far as electron-pair repulsion is concerned double bonds andtriple bonds can be treated like single bonds This approximation is
good for qualitative purposes However you should realize that in
ldquo rdquo
because there are two or three bonds between two atoms theelectron density occupies more space
b) If a molecule has two or more resonance structures we can apply
the VSEPR model to any one of them Formal charges are usually
not shown
c) The order of repulsion strength of lone pair and bond pair are
lone-pair amp lone-pair electrons repulsion are the strongest
followed by lone-pair amp bond-pair electrons repulsion while
bond-pair amp bond-pair electrons repulsion is the weakest
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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32 Covalent bondbull Studies of covalent bond was widely developed ever since
Lewis suggested that a chemical bond exist in a hydrogen gas
occur by sharing en electron between two hydrogen atoms
bull Electron pair that connect the 2 hydrogen atoms is called
covalent bond a bond in which two electrons are shared by
wo a oms an e e ec ron pa r a on e ween e wohydrogen atoms is also called as bonding pair electrons
bull In a covalent bond each electron in a shared pair is attracted to
the nuclei of both atoms This attraction holds the two atoms in
H2 together and is responsible for the formation of covalentbonds in other molecules
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2 A Lewis structure is a representation of covalent bonding
in which shared electron pairs are shown either as lines
or as pairs of dots between two atoms and lone pairsare shown as pairs of dots on individual atoms Only
valence electrons are shown in a Lewis structure
a) Consider the fluorine molecule F2 The electronconfiguration of F is 1s22s22p5 The 1s electrons are in
inner shell which is nearest to the nucleus For this reason
the do not artici ate in bond formation Thus each F
atom has seven valence electrons (2s22p5) Thereforeeach fluorine atom needed one electron to achieve octet
configuration (ns2np6)
F 1 s22s22p5 F 1 s22s22p5 1s2 2s2 2p6 1s2 2s2 2p6
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bull Oxygen atom has electronic configuration of 1s22s22p4 To
achieve stable octet configuration (ns2np6) each oxygen
atom need 2 electrons Hence when 2 oxygen atomsinteract they shared two electrons in between each other as
described in diagram below
O 1s2
2s2
2p4
O 1s2
2s2
2p4
1s2
2s2
2p6
1s2
2s2
2p6
bull From the structure of oxygen molecule formed each oxygen
atom shared two electrons from each other to form a double
bond in order to achieve octet configuration among each
oxygen atom
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bull Nitrogen has electronic configuration of 1s22s22p3 and
required 3 electrons to achieve octet configuration (ns2np6)
In this case each nitrogen atom shared 3 electrons fromeach of its atom to form triple bond
N 1 s22s22p3 N 1 s22s22p3 1s2 2s2 2p6 1s2 2s2 2p6
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Water H2O
H 1 s1 O 2 s22p4
Carbon dioxide CO2
C 2 s22p2 O 2 s22p4
Ammonia NH3
H 1 s1 N 2 s22p3Ethene C2H4
H 1 s1 C 2 s22p2
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Hydrogen cyanide HCN
H 1 s1 N 2 s22p3 C 2 s22p2
Ethanoic acid CH3COOH
H 1 s1 O 2 s22p4 C 2 s22p2
Tetrachloromethane CCl4C 2 s22p2 Cl 3s23p5
Ethyne C2H2
H 1 s1 C 2 s22p2
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a) Note that in ethene hydrogen cyanide and ethanoic acid all
the valence electrons are used in bonding there are no lone
pairs on the carbon atoms In fact most of the stablemolecules containing carbon do not have lone pairs on the
carbon atoms
b) Multiple bonds (double bond or triple bond) are shorter than single covalent bonds Bond length is defined as the
distance between the nuclei of two covalently bonded atoms
nitrogen triple bonds are shorter than double bonds whichin turn are shorter than single bonds The shorter multiple
bonds are also more stable than single bonds
c) Covalent bond not only exist in neutral molecule but also insome molecular ions Table below shows a few example of
molecular ions which have covalent bonds in its molecule
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Carbonate ion CO32-
C 2 s22p2 O 2 s22p4
Cyanide ion CN-
N 2 s22p3 C 2 s22p2
u p a e on 4-
S 3 s23p4 O 2 s22p4
ra e on 3-
N 2 s22p3 O 2 s22p4
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321 Exception of Octet Rules
1 From the Lewis structure sketched for sulphate ion SO42- we can
see that the center metal atom (sulphur) has more than 8electrons For this the molecule is described as molecule that can
expanded octet These molecules that have more than 8
electrons are located at Period 3 and below as these center
atoms have empty d-orbital available to expand the number ofelectrons positioned in the center atom Those center atom from
Period 2 such as C N O and F can only allocate a maximum
Example Both phosphorous and nitrogen are elements from Group15 with the valence electron of ns2np3 They can both form NCl3and PCl3 respectively when react with limited amount of chlorine
however under excess chlorine only PCl 5 can be formed but
not NCl 5 Explain the statement bolded
Solution This is due to phosphorous which is from period
3 have empty d-orbital to expand the octet However
nitrogen which is from Period 2 do not have emptyorbital and can only allocate 8 electrons in its shell
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Phosphorous pentachloride PCl5P 3 s23p3 Cl 3s23p5
Sulphur hexafluoride SF6
S 3 s23p4 F 2 s22p5
Bromine pentachloride BrCl5Cl 3s23p5 Br 4s24p5
Xenon tetrafluoride XeF4
Xe 5s25p6 F 2 s22p5
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2 There are also some stable covalent compounds which have
less than 8 electrons in their center atom (incomplete octet)
The center atoms are usually metals with great number ofvalence electrons with small atomic radius such as
beryllium boron and aluminium
BeCl2Be 2s2 Cl 3s23p5 BF3B 2 s22p1 F 2 s22p5 AlCl3Al 3s23p1 F 3 s23p5
These compound possessed stability due to their short bondlength between center atom with surround atom Furthermore
they can form resonance structure between the center atom
and surrounding atoms
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3 There are also some molecules which contain an odd
number of electrons Among the most common compounds
are nitrogen monoxide (NO) and nitrogen dioxide (NO2)
Odd-electron molecules are sometimes called radicals
Nitrogen monoxide NO
N 2 s22p3 O 2 s22p4
Nitrogen dioxide NO2
N 2 s22p3 O 2 s22p4
any ra ca s are g y reac ve e reason s t at t ere
is a tendency for the unpaired electron to form a
covalent bond with an unpaired electron on another
molecule For example when two nitrogen dioxide
molecules collide they form dinitrogen tetroxide
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322 Dative bond
1 As shown in the Lewis structure of nitrate ion one of the N-Obond is drawn as rarr The bond rarr placed is called as dative
bond (also known as coordinative bond) where dative
bond is defined as a covalent bond in which one of the atoms
donate the lone pair electrons available Although theproperties of a coordinate covalent bond do not differ from
those of a normal covalent bond because all electrons are
a e no ma er w a e r source
2 Dative bond is usually applied for these few circumstances
below
a) To assist atom molecule ion that not yet achieved octet
configuration Making use of atom which has lone pair
electrons to those which are lack of electrons
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Sulphur dioxide SO2Sulphur trioxide SO3 Carbon monoxide CO
Ozone molecule O3 Water with hydrogen ion
S
OO
O
Ammonia with hydrogen ion Ammonia with boron trifluoride
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b) Formation of dimer - In order for some compounds which
have incomplete octet to achieve stability they tend to form
dimer or polymer among themselves by using dative bondTwo of the most common examples are aluminium trichloride
and beryllium dichloride
Monomer of AlCl3Dimer of aluminium chloride Al
2
Cl6
Monomer of BeCl2 Dimer of Be2Cl4 Polymer of (BeCl2)n
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c) Formation of coordination compounds - Coordination
compounds are substances that contain at least one complex
ion a species consisting of a central metal cation (either atransition metal or a main-group metal) that is bonded to
molecules andor anions called ligands via dative
(coordinative) bond
hexaaquacopper (II) ion [Cu(H2O)6]2+ tetraamminenickel (II) ion [Ni(NH3)4]
2+
Hexacyanoferrate (III) ion [Fe(CN)6]3- Trioxalatocobaltate(III)ion [Co(C2O4)3]
3-
3 2 3 Hybridisation Theory
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323 Hybridisation Theory
Valence bond theory
bull The basic principle of valence bond theory is that a covalentbond is formed when orbitals of two atoms overlap and the
overlapped region which is between the nuclei is occupied by
a pair of electrons The central themes of valence bond theory
derive from this principle
ndash Opposing spins of the electron pair ~ Stated in Paulis
orbitals has a maximum capacity of two electrons that musthave opposite spins For example when a covalent bond is
formed in molecule of hydrogen H2 the two 1s electrons of
two H atoms occupy the overlapping 1s orbitals and have
opposite spins
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bull Maximum overlap of bonding orbitals ~ The bond strength
depends on attraction of the nuclei for the shared electrons
so the greater the orbital overlap the stronger (morestable) the bond The extent of overlap depends on the
shapes and directions of the orbitals An s orbital is spherical
but p and d orbitals have more electron density in one
direction than in another Thus whenever possible a bond
involving p or d orbitals will be oriented in the direction that
maximizes overlap For example in hydrogen fluoride (HF)
bond the 1s orbital of H overlaps the half-filled 2p orbital of Falong the long axis of that orbital Any other direction would
result in less overlap and thus a weaker bond
H b idi ti f t i bit l T t f th
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bull Hybridisation of atomic orbitals ~ To account for the
bonding in simple diatomic molecules like HF we picture the
direct overlap of s and p orbitals of isolated atoms But howcan we account for the shapes of so many molecules and
polyatomic ions through the overlap of spherical s orbitals
dumbbell-shaped p orbitals and cloverleaf-shaped d orbitals
Linus Pauling proposed that the valence atomic orbitals in the
molecule are different from those in the isolated atoms The
spatial orientations of these new orbitals lead to more stable
bonds and are consistent with observed molecular shapesThe process of orbital mixing is called hybridisation and
the new atomic orbitals are called hybrid orbitals Two key
points about the number and type of hybrid orbitals are that
i The number of hybrid orbitals obtained equals the number ofatomic orbitals mixed
ii The type of hybrid orbitals obtained varies with the types of
atomic orbitals mixed
3 2 3 1 T f h b idi ti
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3231 Type of hybridisation
1 sp3 hybridisation ~ When four electron groups surround
the central atom the center atom involved must prepare fourorbitals with equal energies to overlap with the four
surrounding electron groups Valence Bond theory uses
hypothetical hybrid orbitals which are atomic orbitals
obtained when two or more non-equivalent orbitals of the
same atom combine in preparation for covalent bond
formation H bridisation is the term a lied to the mixin of
atomic orbitals in an atom (usually a central atom) togenerate a set of hybrid orbitals In the case of sp3 We can
generate four equivalent hybrid orbitals from the center
atom by mixing the s orbital and the three p orbitals
Explanation Energy level diagram Diagram of orbitals
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Explanation Energy level diagram Diagram of orbitals
a) Ground state
Carbon which act as thecenter atom has the valence
electron of 2s22p2
b) Excited state One of the electron from 2s
is promote to 2p orbitals -
equa energy eve
c) Hybridised state
One orbital of 2s and three
orbitals of 2p combined(hybrid) and rearrange
themselves to the shape
and orientation of a
tetrahedral shape
Molecular shape Tetrahedral
Angle between bond pair 10950
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2 sp2 hybridisation ~ When three electron groups surround
the central atom the center atom involved must prepare three
orbitals with equal energies to overlap with the three
surrounding electron groups In sp2 hybridisation three
equivalent (in terms of energy level) from the center atom by
mixing the one s orbital and the two p orbitals Using boron
trifluoride (BF3) as example sp2
hybridisation is explained
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3 sp hybridisation ~ When two electron groups surround
the central atom we observe a linear shape which
means that the bonding orbitals must have a linear
orientation VB theory explains this by proposing thatmixing two nonequivalent orbitals of a central atom one
s and one p gives rise to two equivalent sp hybrid
orbitals that lie 1800 apart
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Beryllium which act as the
center atom has the valence
electron of 2s2
b) Excited state
One of the electron from 2s ispromote to 2p orbitals - equal
energy level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of a
linear shape
Molecular shape Linear
Angle between bond pair 1800
4 The concept of hybridisation is also useful to explain molecules with
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p y p
doubletriple bonds By using the concept of the direct overlapping
orbitals and side-touch lapping orbitals the formation of multiple
bonds in ethene C2H4 and ethyne C2H2 are described
a) Ethene C2H4 - Hybridisation take place for both carbon atoms in
ethene molecule is sp2 hybridisation
bull When both hybridised C is bonded together one of the hybridisedorbital overlapped directly between each other while the other two
hybridised orbitals overlapped directly with hydrogen atoms
z
atoms form a side-touch bond between each other and form anotherbond as shown in the diagram below
bull From the diagram with C=C there are two types of bond A sigma-
bond (σ-bond) is covalent bonds formed by orbitals overlapping end-
to-end with the electron density concentrated between the nuclei ofthe bonding atoms while the second type is called a pi bond (π-
bond) which is defined as a covalent bond formed by sideways
overlapping orbitals with electron density concentrated above and
below the plane of the nuclei of the bonding atoms
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital of 2s and two
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of atrigonal planar shape
Note that on unhybridised pz
orbital an electron is presence
Molecular shape Trigonal planar Angle between bond pair 1200
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b) Ethyne C H has the Lewis structure of H CequivC H which the
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b) Ethyne C2H2 has the Lewis structure of HminusCequivCminusH which the
bonding can be explain using sp hybridisation Table below
described how the hybridisation take place on each carbonatom and how the formation of triple bond occur
c) From the diagram formation of -CequivC- is due to the formation
of one sigma-bond by direct overlapping of one of the twohybrid orbitals on each C while the other two pi-bond bonds
are formed as a result of side-lapping of the each two
y z
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital from 2s and 2p
orbitals combined (hybrid) and
rearrange themselves to the
shape and orientation of a
linear shapeNote that on unhybridised py amp
pz orbitals electrons are
presence to form two
π-bonds Molecular shape Linear Angle between bond pair 1800
sp
pz py
5 Other examples of applications in valence bond theory includes the
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formation of nitrogen molecule N2 and hydrogen cyanide HCN
molecule
a) Nitrogen gas is Earths most abundant gas as it cover 78 of the
content of our air Nitrogen molecule is an inert gas thanks to its
short covalent bond and also its strong triple bond Therefore
a lot of heats are required to break the chemical bond of nitrogenbefore it can be applied in industries Using valence bond theory
the bonding of nitrogen molecule is explained in the diagram
b) So when the two hybridised nitrogen atom interacting amongeach other they hence formed a linear shape and two pi-bonds
(πminusbonds) are formed as a result of side-lapping of unhybridised
py and pz orbital respectively
N N
ExplanationEnergy level
diagramDiagram of orbitals
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diagram
a) Ground state
Both nitrogen whichact as the center atom
has the valence
electron of 2s22p3
b) Excited state One of the electron
from 2s is promote to 2p
orbitals - e ual ener
level
c) Hybridised state
One orbital of 2s and
one orbitals of 2p
combined (hybrid) and
rearrange themselves
to the shape and
orientation of a linear
shapeMolecular shape Linear
Angle between bond pair 1800
b) Whereas for hydrogen cyanide HCN both carbon atom and
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) y g y
nitrogen atom undergoes sp hybridisation in order to form a
linear structureExplanation
Energy level diagram for
carbon atom C
Energy level diagram for
nitrogen atom N
a) Ground state
Valence electronC 2s22p2
N 2s22p3
b) Excited state
One of the electron from
2s is promote to 2p
orbitals - equal energy
level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined
(hybrid) and form sp
hybridisation
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N
σσσσ
π
σσσσ
π
6 However there are a few limitation on valence bond theory such
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as when explaining the effect angle of bond-pair and bond-pair
electrons when there isare presence of lone pair electron in thecenter atom and also the difference of electronegativity
a) Bonding in ammonia NH3 and water H2O ~ Both NH3 and
H2O undergoes an arrangement similar to sp3 hybridisation
similar to that of C in methane Table below shows thehybridisation occur for nitrogen in ammonia and oxygen in water
Energy level diagram Energy level diagram
for nitrogen atom N for oxygen atom O
a) Ground state
Valence electron
N 2s22p3
O 2s2
2p4
b) Excited state
One of the electron
from 2s is promote to
2p orbitals
c) Hybridised state
One orbital of 2s and
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One orbital of 2s and
three orbitals from 2p
combined (hybrid) andrearrange themselves to
the shape and
orientation similar to that
of tetrahedral shape
Similar to arrangement in
tetrahedral
Similar to arrangement in
tetrahedral
of ammonia NH3 andwater H2O
Shape and angleShape trigonal pyramidal
Angle 1070
Shape bent
Angle 10450
Number of bond pair amp
lone pair electrons
Bond pair electron 3
Lone pair electron 1
Bond pair electron 2
Lone pair electron 2
bull The angle of bond pair - bond pair electrons in ammonia and0 0
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water are 1070 and 10450 respectively which is lesser than in
methane molecule (10950
) This can be explained by the factof the presence of lone-pair electrons in both ammonia and
water Since the lone pair - lone pair electron repulsion is
stronger than lone pair - bond pair electron repulsion than
bond pair - bond pair electron repulsion it is expected that therepulsion between lone pair - bond pair electrons in ammonia
is stronger than bond pair - bond pair electrons repulsion
hence caused the angle to squeeze to a smaller angle As
for water since there is a presence of lone pair - lone pair
electrons repulsion it results the bond pair - bond pair
electron repulsion to be much smaller hence caused the
angle to squeeze to a smaller angle Though valence bondtheory does not actually explain the hybridisation especially
when it relates to the repulsion occur involving lone pair
electron another theories shall be applied to study such
effect All these shall be discussed further in VSEPR theory
b) Effect of electronegativity and bonding angle between bond pair
- bond pair electrons ~
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- bond pair electrons
i As discussed earlier in the bonding of water (H2
O) the
molecular shape and angle is described in the diagram below
Sul hur which is also an element from Grou 16 formed h dro en
sulphide H2S when sulphur react with hydrogen However unlike
water the bonding angles is much smaller compare to water This is
due to the difference of electronegativity and also the bond length
between O and S in molecule Since O is more electronegative
compare to S the O-H bond is pulled closer toward O Furthermorethe bond length of O-H is shorter compare to S-H As a result the
bonding pair-bonding pair electrons repulse greater with each other in
H-O-H and caused the angle become greater
ii Another example is between NH3 and PH3 The orbital
f
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diagram for both NH3 and PH3 are described below Both
nitrogen N and phosphorous P are from the same groupwhich is Group 15
Similar to the above case P in phosphine is less
electronegative than N in ammonia and bond length of N-H
is shorter than P-H As a result H is pull closer to N and
repulsion between bonding of H-N-H is greater compare toH-P-H and caused the angle between H-N-H is greater
compare to H-P-H
324 Valence Shell Electron Pair Repulsion (VSEPR) Theory
1 M l l t i th th di i l t f t
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1 Molecular geometry is the three-dimensional arrangement of atoms
in a molecule A moleculersquos geometry affects its physical andchemical properties such as melting point boiling point density
and the types of reactions it undergoes The basic concept of
VSEPR theory is based on the three general rules below
a) As far as electron-pair repulsion is concerned double bonds andtriple bonds can be treated like single bonds This approximation is
good for qualitative purposes However you should realize that in
ldquo rdquo
because there are two or three bonds between two atoms theelectron density occupies more space
b) If a molecule has two or more resonance structures we can apply
the VSEPR model to any one of them Formal charges are usually
not shown
c) The order of repulsion strength of lone pair and bond pair are
lone-pair amp lone-pair electrons repulsion are the strongest
followed by lone-pair amp bond-pair electrons repulsion while
bond-pair amp bond-pair electrons repulsion is the weakest
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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32 Covalent bondbull Studies of covalent bond was widely developed ever since
Lewis suggested that a chemical bond exist in a hydrogen gas
occur by sharing en electron between two hydrogen atoms
bull Electron pair that connect the 2 hydrogen atoms is called
covalent bond a bond in which two electrons are shared by
wo a oms an e e ec ron pa r a on e ween e wohydrogen atoms is also called as bonding pair electrons
bull In a covalent bond each electron in a shared pair is attracted to
the nuclei of both atoms This attraction holds the two atoms in
H2 together and is responsible for the formation of covalentbonds in other molecules
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2 A Lewis structure is a representation of covalent bonding
in which shared electron pairs are shown either as lines
or as pairs of dots between two atoms and lone pairsare shown as pairs of dots on individual atoms Only
valence electrons are shown in a Lewis structure
a) Consider the fluorine molecule F2 The electronconfiguration of F is 1s22s22p5 The 1s electrons are in
inner shell which is nearest to the nucleus For this reason
the do not artici ate in bond formation Thus each F
atom has seven valence electrons (2s22p5) Thereforeeach fluorine atom needed one electron to achieve octet
configuration (ns2np6)
F 1 s22s22p5 F 1 s22s22p5 1s2 2s2 2p6 1s2 2s2 2p6
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bull Oxygen atom has electronic configuration of 1s22s22p4 To
achieve stable octet configuration (ns2np6) each oxygen
atom need 2 electrons Hence when 2 oxygen atomsinteract they shared two electrons in between each other as
described in diagram below
O 1s2
2s2
2p4
O 1s2
2s2
2p4
1s2
2s2
2p6
1s2
2s2
2p6
bull From the structure of oxygen molecule formed each oxygen
atom shared two electrons from each other to form a double
bond in order to achieve octet configuration among each
oxygen atom
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bull Nitrogen has electronic configuration of 1s22s22p3 and
required 3 electrons to achieve octet configuration (ns2np6)
In this case each nitrogen atom shared 3 electrons fromeach of its atom to form triple bond
N 1 s22s22p3 N 1 s22s22p3 1s2 2s2 2p6 1s2 2s2 2p6
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Water H2O
H 1 s1 O 2 s22p4
Carbon dioxide CO2
C 2 s22p2 O 2 s22p4
Ammonia NH3
H 1 s1 N 2 s22p3Ethene C2H4
H 1 s1 C 2 s22p2
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Hydrogen cyanide HCN
H 1 s1 N 2 s22p3 C 2 s22p2
Ethanoic acid CH3COOH
H 1 s1 O 2 s22p4 C 2 s22p2
Tetrachloromethane CCl4C 2 s22p2 Cl 3s23p5
Ethyne C2H2
H 1 s1 C 2 s22p2
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a) Note that in ethene hydrogen cyanide and ethanoic acid all
the valence electrons are used in bonding there are no lone
pairs on the carbon atoms In fact most of the stablemolecules containing carbon do not have lone pairs on the
carbon atoms
b) Multiple bonds (double bond or triple bond) are shorter than single covalent bonds Bond length is defined as the
distance between the nuclei of two covalently bonded atoms
nitrogen triple bonds are shorter than double bonds whichin turn are shorter than single bonds The shorter multiple
bonds are also more stable than single bonds
c) Covalent bond not only exist in neutral molecule but also insome molecular ions Table below shows a few example of
molecular ions which have covalent bonds in its molecule
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Carbonate ion CO32-
C 2 s22p2 O 2 s22p4
Cyanide ion CN-
N 2 s22p3 C 2 s22p2
u p a e on 4-
S 3 s23p4 O 2 s22p4
ra e on 3-
N 2 s22p3 O 2 s22p4
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321 Exception of Octet Rules
1 From the Lewis structure sketched for sulphate ion SO42- we can
see that the center metal atom (sulphur) has more than 8electrons For this the molecule is described as molecule that can
expanded octet These molecules that have more than 8
electrons are located at Period 3 and below as these center
atoms have empty d-orbital available to expand the number ofelectrons positioned in the center atom Those center atom from
Period 2 such as C N O and F can only allocate a maximum
Example Both phosphorous and nitrogen are elements from Group15 with the valence electron of ns2np3 They can both form NCl3and PCl3 respectively when react with limited amount of chlorine
however under excess chlorine only PCl 5 can be formed but
not NCl 5 Explain the statement bolded
Solution This is due to phosphorous which is from period
3 have empty d-orbital to expand the octet However
nitrogen which is from Period 2 do not have emptyorbital and can only allocate 8 electrons in its shell
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Phosphorous pentachloride PCl5P 3 s23p3 Cl 3s23p5
Sulphur hexafluoride SF6
S 3 s23p4 F 2 s22p5
Bromine pentachloride BrCl5Cl 3s23p5 Br 4s24p5
Xenon tetrafluoride XeF4
Xe 5s25p6 F 2 s22p5
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2 There are also some stable covalent compounds which have
less than 8 electrons in their center atom (incomplete octet)
The center atoms are usually metals with great number ofvalence electrons with small atomic radius such as
beryllium boron and aluminium
BeCl2Be 2s2 Cl 3s23p5 BF3B 2 s22p1 F 2 s22p5 AlCl3Al 3s23p1 F 3 s23p5
These compound possessed stability due to their short bondlength between center atom with surround atom Furthermore
they can form resonance structure between the center atom
and surrounding atoms
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3 There are also some molecules which contain an odd
number of electrons Among the most common compounds
are nitrogen monoxide (NO) and nitrogen dioxide (NO2)
Odd-electron molecules are sometimes called radicals
Nitrogen monoxide NO
N 2 s22p3 O 2 s22p4
Nitrogen dioxide NO2
N 2 s22p3 O 2 s22p4
any ra ca s are g y reac ve e reason s t at t ere
is a tendency for the unpaired electron to form a
covalent bond with an unpaired electron on another
molecule For example when two nitrogen dioxide
molecules collide they form dinitrogen tetroxide
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322 Dative bond
1 As shown in the Lewis structure of nitrate ion one of the N-Obond is drawn as rarr The bond rarr placed is called as dative
bond (also known as coordinative bond) where dative
bond is defined as a covalent bond in which one of the atoms
donate the lone pair electrons available Although theproperties of a coordinate covalent bond do not differ from
those of a normal covalent bond because all electrons are
a e no ma er w a e r source
2 Dative bond is usually applied for these few circumstances
below
a) To assist atom molecule ion that not yet achieved octet
configuration Making use of atom which has lone pair
electrons to those which are lack of electrons
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Sulphur dioxide SO2Sulphur trioxide SO3 Carbon monoxide CO
Ozone molecule O3 Water with hydrogen ion
S
OO
O
Ammonia with hydrogen ion Ammonia with boron trifluoride
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b) Formation of dimer - In order for some compounds which
have incomplete octet to achieve stability they tend to form
dimer or polymer among themselves by using dative bondTwo of the most common examples are aluminium trichloride
and beryllium dichloride
Monomer of AlCl3Dimer of aluminium chloride Al
2
Cl6
Monomer of BeCl2 Dimer of Be2Cl4 Polymer of (BeCl2)n
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c) Formation of coordination compounds - Coordination
compounds are substances that contain at least one complex
ion a species consisting of a central metal cation (either atransition metal or a main-group metal) that is bonded to
molecules andor anions called ligands via dative
(coordinative) bond
hexaaquacopper (II) ion [Cu(H2O)6]2+ tetraamminenickel (II) ion [Ni(NH3)4]
2+
Hexacyanoferrate (III) ion [Fe(CN)6]3- Trioxalatocobaltate(III)ion [Co(C2O4)3]
3-
3 2 3 Hybridisation Theory
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323 Hybridisation Theory
Valence bond theory
bull The basic principle of valence bond theory is that a covalentbond is formed when orbitals of two atoms overlap and the
overlapped region which is between the nuclei is occupied by
a pair of electrons The central themes of valence bond theory
derive from this principle
ndash Opposing spins of the electron pair ~ Stated in Paulis
orbitals has a maximum capacity of two electrons that musthave opposite spins For example when a covalent bond is
formed in molecule of hydrogen H2 the two 1s electrons of
two H atoms occupy the overlapping 1s orbitals and have
opposite spins
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bull Maximum overlap of bonding orbitals ~ The bond strength
depends on attraction of the nuclei for the shared electrons
so the greater the orbital overlap the stronger (morestable) the bond The extent of overlap depends on the
shapes and directions of the orbitals An s orbital is spherical
but p and d orbitals have more electron density in one
direction than in another Thus whenever possible a bond
involving p or d orbitals will be oriented in the direction that
maximizes overlap For example in hydrogen fluoride (HF)
bond the 1s orbital of H overlaps the half-filled 2p orbital of Falong the long axis of that orbital Any other direction would
result in less overlap and thus a weaker bond
H b idi ti f t i bit l T t f th
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bull Hybridisation of atomic orbitals ~ To account for the
bonding in simple diatomic molecules like HF we picture the
direct overlap of s and p orbitals of isolated atoms But howcan we account for the shapes of so many molecules and
polyatomic ions through the overlap of spherical s orbitals
dumbbell-shaped p orbitals and cloverleaf-shaped d orbitals
Linus Pauling proposed that the valence atomic orbitals in the
molecule are different from those in the isolated atoms The
spatial orientations of these new orbitals lead to more stable
bonds and are consistent with observed molecular shapesThe process of orbital mixing is called hybridisation and
the new atomic orbitals are called hybrid orbitals Two key
points about the number and type of hybrid orbitals are that
i The number of hybrid orbitals obtained equals the number ofatomic orbitals mixed
ii The type of hybrid orbitals obtained varies with the types of
atomic orbitals mixed
3 2 3 1 T f h b idi ti
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3231 Type of hybridisation
1 sp3 hybridisation ~ When four electron groups surround
the central atom the center atom involved must prepare fourorbitals with equal energies to overlap with the four
surrounding electron groups Valence Bond theory uses
hypothetical hybrid orbitals which are atomic orbitals
obtained when two or more non-equivalent orbitals of the
same atom combine in preparation for covalent bond
formation H bridisation is the term a lied to the mixin of
atomic orbitals in an atom (usually a central atom) togenerate a set of hybrid orbitals In the case of sp3 We can
generate four equivalent hybrid orbitals from the center
atom by mixing the s orbital and the three p orbitals
Explanation Energy level diagram Diagram of orbitals
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Explanation Energy level diagram Diagram of orbitals
a) Ground state
Carbon which act as thecenter atom has the valence
electron of 2s22p2
b) Excited state One of the electron from 2s
is promote to 2p orbitals -
equa energy eve
c) Hybridised state
One orbital of 2s and three
orbitals of 2p combined(hybrid) and rearrange
themselves to the shape
and orientation of a
tetrahedral shape
Molecular shape Tetrahedral
Angle between bond pair 10950
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2 sp2 hybridisation ~ When three electron groups surround
the central atom the center atom involved must prepare three
orbitals with equal energies to overlap with the three
surrounding electron groups In sp2 hybridisation three
equivalent (in terms of energy level) from the center atom by
mixing the one s orbital and the two p orbitals Using boron
trifluoride (BF3) as example sp2
hybridisation is explained
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3 sp hybridisation ~ When two electron groups surround
the central atom we observe a linear shape which
means that the bonding orbitals must have a linear
orientation VB theory explains this by proposing thatmixing two nonequivalent orbitals of a central atom one
s and one p gives rise to two equivalent sp hybrid
orbitals that lie 1800 apart
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Beryllium which act as the
center atom has the valence
electron of 2s2
b) Excited state
One of the electron from 2s ispromote to 2p orbitals - equal
energy level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of a
linear shape
Molecular shape Linear
Angle between bond pair 1800
4 The concept of hybridisation is also useful to explain molecules with
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p y p
doubletriple bonds By using the concept of the direct overlapping
orbitals and side-touch lapping orbitals the formation of multiple
bonds in ethene C2H4 and ethyne C2H2 are described
a) Ethene C2H4 - Hybridisation take place for both carbon atoms in
ethene molecule is sp2 hybridisation
bull When both hybridised C is bonded together one of the hybridisedorbital overlapped directly between each other while the other two
hybridised orbitals overlapped directly with hydrogen atoms
z
atoms form a side-touch bond between each other and form anotherbond as shown in the diagram below
bull From the diagram with C=C there are two types of bond A sigma-
bond (σ-bond) is covalent bonds formed by orbitals overlapping end-
to-end with the electron density concentrated between the nuclei ofthe bonding atoms while the second type is called a pi bond (π-
bond) which is defined as a covalent bond formed by sideways
overlapping orbitals with electron density concentrated above and
below the plane of the nuclei of the bonding atoms
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital of 2s and two
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of atrigonal planar shape
Note that on unhybridised pz
orbital an electron is presence
Molecular shape Trigonal planar Angle between bond pair 1200
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b) Ethyne C H has the Lewis structure of H CequivC H which the
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b) Ethyne C2H2 has the Lewis structure of HminusCequivCminusH which the
bonding can be explain using sp hybridisation Table below
described how the hybridisation take place on each carbonatom and how the formation of triple bond occur
c) From the diagram formation of -CequivC- is due to the formation
of one sigma-bond by direct overlapping of one of the twohybrid orbitals on each C while the other two pi-bond bonds
are formed as a result of side-lapping of the each two
y z
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital from 2s and 2p
orbitals combined (hybrid) and
rearrange themselves to the
shape and orientation of a
linear shapeNote that on unhybridised py amp
pz orbitals electrons are
presence to form two
π-bonds Molecular shape Linear Angle between bond pair 1800
sp
pz py
5 Other examples of applications in valence bond theory includes the
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formation of nitrogen molecule N2 and hydrogen cyanide HCN
molecule
a) Nitrogen gas is Earths most abundant gas as it cover 78 of the
content of our air Nitrogen molecule is an inert gas thanks to its
short covalent bond and also its strong triple bond Therefore
a lot of heats are required to break the chemical bond of nitrogenbefore it can be applied in industries Using valence bond theory
the bonding of nitrogen molecule is explained in the diagram
b) So when the two hybridised nitrogen atom interacting amongeach other they hence formed a linear shape and two pi-bonds
(πminusbonds) are formed as a result of side-lapping of unhybridised
py and pz orbital respectively
N N
ExplanationEnergy level
diagramDiagram of orbitals
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diagram
a) Ground state
Both nitrogen whichact as the center atom
has the valence
electron of 2s22p3
b) Excited state One of the electron
from 2s is promote to 2p
orbitals - e ual ener
level
c) Hybridised state
One orbital of 2s and
one orbitals of 2p
combined (hybrid) and
rearrange themselves
to the shape and
orientation of a linear
shapeMolecular shape Linear
Angle between bond pair 1800
b) Whereas for hydrogen cyanide HCN both carbon atom and
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) y g y
nitrogen atom undergoes sp hybridisation in order to form a
linear structureExplanation
Energy level diagram for
carbon atom C
Energy level diagram for
nitrogen atom N
a) Ground state
Valence electronC 2s22p2
N 2s22p3
b) Excited state
One of the electron from
2s is promote to 2p
orbitals - equal energy
level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined
(hybrid) and form sp
hybridisation
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N
σσσσ
π
σσσσ
π
6 However there are a few limitation on valence bond theory such
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as when explaining the effect angle of bond-pair and bond-pair
electrons when there isare presence of lone pair electron in thecenter atom and also the difference of electronegativity
a) Bonding in ammonia NH3 and water H2O ~ Both NH3 and
H2O undergoes an arrangement similar to sp3 hybridisation
similar to that of C in methane Table below shows thehybridisation occur for nitrogen in ammonia and oxygen in water
Energy level diagram Energy level diagram
for nitrogen atom N for oxygen atom O
a) Ground state
Valence electron
N 2s22p3
O 2s2
2p4
b) Excited state
One of the electron
from 2s is promote to
2p orbitals
c) Hybridised state
One orbital of 2s and
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One orbital of 2s and
three orbitals from 2p
combined (hybrid) andrearrange themselves to
the shape and
orientation similar to that
of tetrahedral shape
Similar to arrangement in
tetrahedral
Similar to arrangement in
tetrahedral
of ammonia NH3 andwater H2O
Shape and angleShape trigonal pyramidal
Angle 1070
Shape bent
Angle 10450
Number of bond pair amp
lone pair electrons
Bond pair electron 3
Lone pair electron 1
Bond pair electron 2
Lone pair electron 2
bull The angle of bond pair - bond pair electrons in ammonia and0 0
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water are 1070 and 10450 respectively which is lesser than in
methane molecule (10950
) This can be explained by the factof the presence of lone-pair electrons in both ammonia and
water Since the lone pair - lone pair electron repulsion is
stronger than lone pair - bond pair electron repulsion than
bond pair - bond pair electron repulsion it is expected that therepulsion between lone pair - bond pair electrons in ammonia
is stronger than bond pair - bond pair electrons repulsion
hence caused the angle to squeeze to a smaller angle As
for water since there is a presence of lone pair - lone pair
electrons repulsion it results the bond pair - bond pair
electron repulsion to be much smaller hence caused the
angle to squeeze to a smaller angle Though valence bondtheory does not actually explain the hybridisation especially
when it relates to the repulsion occur involving lone pair
electron another theories shall be applied to study such
effect All these shall be discussed further in VSEPR theory
b) Effect of electronegativity and bonding angle between bond pair
- bond pair electrons ~
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- bond pair electrons
i As discussed earlier in the bonding of water (H2
O) the
molecular shape and angle is described in the diagram below
Sul hur which is also an element from Grou 16 formed h dro en
sulphide H2S when sulphur react with hydrogen However unlike
water the bonding angles is much smaller compare to water This is
due to the difference of electronegativity and also the bond length
between O and S in molecule Since O is more electronegative
compare to S the O-H bond is pulled closer toward O Furthermorethe bond length of O-H is shorter compare to S-H As a result the
bonding pair-bonding pair electrons repulse greater with each other in
H-O-H and caused the angle become greater
ii Another example is between NH3 and PH3 The orbital
f
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diagram for both NH3 and PH3 are described below Both
nitrogen N and phosphorous P are from the same groupwhich is Group 15
Similar to the above case P in phosphine is less
electronegative than N in ammonia and bond length of N-H
is shorter than P-H As a result H is pull closer to N and
repulsion between bonding of H-N-H is greater compare toH-P-H and caused the angle between H-N-H is greater
compare to H-P-H
324 Valence Shell Electron Pair Repulsion (VSEPR) Theory
1 M l l t i th th di i l t f t
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1 Molecular geometry is the three-dimensional arrangement of atoms
in a molecule A moleculersquos geometry affects its physical andchemical properties such as melting point boiling point density
and the types of reactions it undergoes The basic concept of
VSEPR theory is based on the three general rules below
a) As far as electron-pair repulsion is concerned double bonds andtriple bonds can be treated like single bonds This approximation is
good for qualitative purposes However you should realize that in
ldquo rdquo
because there are two or three bonds between two atoms theelectron density occupies more space
b) If a molecule has two or more resonance structures we can apply
the VSEPR model to any one of them Formal charges are usually
not shown
c) The order of repulsion strength of lone pair and bond pair are
lone-pair amp lone-pair electrons repulsion are the strongest
followed by lone-pair amp bond-pair electrons repulsion while
bond-pair amp bond-pair electrons repulsion is the weakest
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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32 Covalent bondbull Studies of covalent bond was widely developed ever since
Lewis suggested that a chemical bond exist in a hydrogen gas
occur by sharing en electron between two hydrogen atoms
bull Electron pair that connect the 2 hydrogen atoms is called
covalent bond a bond in which two electrons are shared by
wo a oms an e e ec ron pa r a on e ween e wohydrogen atoms is also called as bonding pair electrons
bull In a covalent bond each electron in a shared pair is attracted to
the nuclei of both atoms This attraction holds the two atoms in
H2 together and is responsible for the formation of covalentbonds in other molecules
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2 A Lewis structure is a representation of covalent bonding
in which shared electron pairs are shown either as lines
or as pairs of dots between two atoms and lone pairsare shown as pairs of dots on individual atoms Only
valence electrons are shown in a Lewis structure
a) Consider the fluorine molecule F2 The electronconfiguration of F is 1s22s22p5 The 1s electrons are in
inner shell which is nearest to the nucleus For this reason
the do not artici ate in bond formation Thus each F
atom has seven valence electrons (2s22p5) Thereforeeach fluorine atom needed one electron to achieve octet
configuration (ns2np6)
F 1 s22s22p5 F 1 s22s22p5 1s2 2s2 2p6 1s2 2s2 2p6
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bull Oxygen atom has electronic configuration of 1s22s22p4 To
achieve stable octet configuration (ns2np6) each oxygen
atom need 2 electrons Hence when 2 oxygen atomsinteract they shared two electrons in between each other as
described in diagram below
O 1s2
2s2
2p4
O 1s2
2s2
2p4
1s2
2s2
2p6
1s2
2s2
2p6
bull From the structure of oxygen molecule formed each oxygen
atom shared two electrons from each other to form a double
bond in order to achieve octet configuration among each
oxygen atom
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bull Nitrogen has electronic configuration of 1s22s22p3 and
required 3 electrons to achieve octet configuration (ns2np6)
In this case each nitrogen atom shared 3 electrons fromeach of its atom to form triple bond
N 1 s22s22p3 N 1 s22s22p3 1s2 2s2 2p6 1s2 2s2 2p6
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Water H2O
H 1 s1 O 2 s22p4
Carbon dioxide CO2
C 2 s22p2 O 2 s22p4
Ammonia NH3
H 1 s1 N 2 s22p3Ethene C2H4
H 1 s1 C 2 s22p2
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Hydrogen cyanide HCN
H 1 s1 N 2 s22p3 C 2 s22p2
Ethanoic acid CH3COOH
H 1 s1 O 2 s22p4 C 2 s22p2
Tetrachloromethane CCl4C 2 s22p2 Cl 3s23p5
Ethyne C2H2
H 1 s1 C 2 s22p2
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a) Note that in ethene hydrogen cyanide and ethanoic acid all
the valence electrons are used in bonding there are no lone
pairs on the carbon atoms In fact most of the stablemolecules containing carbon do not have lone pairs on the
carbon atoms
b) Multiple bonds (double bond or triple bond) are shorter than single covalent bonds Bond length is defined as the
distance between the nuclei of two covalently bonded atoms
nitrogen triple bonds are shorter than double bonds whichin turn are shorter than single bonds The shorter multiple
bonds are also more stable than single bonds
c) Covalent bond not only exist in neutral molecule but also insome molecular ions Table below shows a few example of
molecular ions which have covalent bonds in its molecule
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Carbonate ion CO32-
C 2 s22p2 O 2 s22p4
Cyanide ion CN-
N 2 s22p3 C 2 s22p2
u p a e on 4-
S 3 s23p4 O 2 s22p4
ra e on 3-
N 2 s22p3 O 2 s22p4
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321 Exception of Octet Rules
1 From the Lewis structure sketched for sulphate ion SO42- we can
see that the center metal atom (sulphur) has more than 8electrons For this the molecule is described as molecule that can
expanded octet These molecules that have more than 8
electrons are located at Period 3 and below as these center
atoms have empty d-orbital available to expand the number ofelectrons positioned in the center atom Those center atom from
Period 2 such as C N O and F can only allocate a maximum
Example Both phosphorous and nitrogen are elements from Group15 with the valence electron of ns2np3 They can both form NCl3and PCl3 respectively when react with limited amount of chlorine
however under excess chlorine only PCl 5 can be formed but
not NCl 5 Explain the statement bolded
Solution This is due to phosphorous which is from period
3 have empty d-orbital to expand the octet However
nitrogen which is from Period 2 do not have emptyorbital and can only allocate 8 electrons in its shell
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Phosphorous pentachloride PCl5P 3 s23p3 Cl 3s23p5
Sulphur hexafluoride SF6
S 3 s23p4 F 2 s22p5
Bromine pentachloride BrCl5Cl 3s23p5 Br 4s24p5
Xenon tetrafluoride XeF4
Xe 5s25p6 F 2 s22p5
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2 There are also some stable covalent compounds which have
less than 8 electrons in their center atom (incomplete octet)
The center atoms are usually metals with great number ofvalence electrons with small atomic radius such as
beryllium boron and aluminium
BeCl2Be 2s2 Cl 3s23p5 BF3B 2 s22p1 F 2 s22p5 AlCl3Al 3s23p1 F 3 s23p5
These compound possessed stability due to their short bondlength between center atom with surround atom Furthermore
they can form resonance structure between the center atom
and surrounding atoms
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3 There are also some molecules which contain an odd
number of electrons Among the most common compounds
are nitrogen monoxide (NO) and nitrogen dioxide (NO2)
Odd-electron molecules are sometimes called radicals
Nitrogen monoxide NO
N 2 s22p3 O 2 s22p4
Nitrogen dioxide NO2
N 2 s22p3 O 2 s22p4
any ra ca s are g y reac ve e reason s t at t ere
is a tendency for the unpaired electron to form a
covalent bond with an unpaired electron on another
molecule For example when two nitrogen dioxide
molecules collide they form dinitrogen tetroxide
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322 Dative bond
1 As shown in the Lewis structure of nitrate ion one of the N-Obond is drawn as rarr The bond rarr placed is called as dative
bond (also known as coordinative bond) where dative
bond is defined as a covalent bond in which one of the atoms
donate the lone pair electrons available Although theproperties of a coordinate covalent bond do not differ from
those of a normal covalent bond because all electrons are
a e no ma er w a e r source
2 Dative bond is usually applied for these few circumstances
below
a) To assist atom molecule ion that not yet achieved octet
configuration Making use of atom which has lone pair
electrons to those which are lack of electrons
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Sulphur dioxide SO2Sulphur trioxide SO3 Carbon monoxide CO
Ozone molecule O3 Water with hydrogen ion
S
OO
O
Ammonia with hydrogen ion Ammonia with boron trifluoride
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b) Formation of dimer - In order for some compounds which
have incomplete octet to achieve stability they tend to form
dimer or polymer among themselves by using dative bondTwo of the most common examples are aluminium trichloride
and beryllium dichloride
Monomer of AlCl3Dimer of aluminium chloride Al
2
Cl6
Monomer of BeCl2 Dimer of Be2Cl4 Polymer of (BeCl2)n
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c) Formation of coordination compounds - Coordination
compounds are substances that contain at least one complex
ion a species consisting of a central metal cation (either atransition metal or a main-group metal) that is bonded to
molecules andor anions called ligands via dative
(coordinative) bond
hexaaquacopper (II) ion [Cu(H2O)6]2+ tetraamminenickel (II) ion [Ni(NH3)4]
2+
Hexacyanoferrate (III) ion [Fe(CN)6]3- Trioxalatocobaltate(III)ion [Co(C2O4)3]
3-
3 2 3 Hybridisation Theory
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323 Hybridisation Theory
Valence bond theory
bull The basic principle of valence bond theory is that a covalentbond is formed when orbitals of two atoms overlap and the
overlapped region which is between the nuclei is occupied by
a pair of electrons The central themes of valence bond theory
derive from this principle
ndash Opposing spins of the electron pair ~ Stated in Paulis
orbitals has a maximum capacity of two electrons that musthave opposite spins For example when a covalent bond is
formed in molecule of hydrogen H2 the two 1s electrons of
two H atoms occupy the overlapping 1s orbitals and have
opposite spins
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bull Maximum overlap of bonding orbitals ~ The bond strength
depends on attraction of the nuclei for the shared electrons
so the greater the orbital overlap the stronger (morestable) the bond The extent of overlap depends on the
shapes and directions of the orbitals An s orbital is spherical
but p and d orbitals have more electron density in one
direction than in another Thus whenever possible a bond
involving p or d orbitals will be oriented in the direction that
maximizes overlap For example in hydrogen fluoride (HF)
bond the 1s orbital of H overlaps the half-filled 2p orbital of Falong the long axis of that orbital Any other direction would
result in less overlap and thus a weaker bond
H b idi ti f t i bit l T t f th
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bull Hybridisation of atomic orbitals ~ To account for the
bonding in simple diatomic molecules like HF we picture the
direct overlap of s and p orbitals of isolated atoms But howcan we account for the shapes of so many molecules and
polyatomic ions through the overlap of spherical s orbitals
dumbbell-shaped p orbitals and cloverleaf-shaped d orbitals
Linus Pauling proposed that the valence atomic orbitals in the
molecule are different from those in the isolated atoms The
spatial orientations of these new orbitals lead to more stable
bonds and are consistent with observed molecular shapesThe process of orbital mixing is called hybridisation and
the new atomic orbitals are called hybrid orbitals Two key
points about the number and type of hybrid orbitals are that
i The number of hybrid orbitals obtained equals the number ofatomic orbitals mixed
ii The type of hybrid orbitals obtained varies with the types of
atomic orbitals mixed
3 2 3 1 T f h b idi ti
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3231 Type of hybridisation
1 sp3 hybridisation ~ When four electron groups surround
the central atom the center atom involved must prepare fourorbitals with equal energies to overlap with the four
surrounding electron groups Valence Bond theory uses
hypothetical hybrid orbitals which are atomic orbitals
obtained when two or more non-equivalent orbitals of the
same atom combine in preparation for covalent bond
formation H bridisation is the term a lied to the mixin of
atomic orbitals in an atom (usually a central atom) togenerate a set of hybrid orbitals In the case of sp3 We can
generate four equivalent hybrid orbitals from the center
atom by mixing the s orbital and the three p orbitals
Explanation Energy level diagram Diagram of orbitals
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Explanation Energy level diagram Diagram of orbitals
a) Ground state
Carbon which act as thecenter atom has the valence
electron of 2s22p2
b) Excited state One of the electron from 2s
is promote to 2p orbitals -
equa energy eve
c) Hybridised state
One orbital of 2s and three
orbitals of 2p combined(hybrid) and rearrange
themselves to the shape
and orientation of a
tetrahedral shape
Molecular shape Tetrahedral
Angle between bond pair 10950
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2 sp2 hybridisation ~ When three electron groups surround
the central atom the center atom involved must prepare three
orbitals with equal energies to overlap with the three
surrounding electron groups In sp2 hybridisation three
equivalent (in terms of energy level) from the center atom by
mixing the one s orbital and the two p orbitals Using boron
trifluoride (BF3) as example sp2
hybridisation is explained
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3 sp hybridisation ~ When two electron groups surround
the central atom we observe a linear shape which
means that the bonding orbitals must have a linear
orientation VB theory explains this by proposing thatmixing two nonequivalent orbitals of a central atom one
s and one p gives rise to two equivalent sp hybrid
orbitals that lie 1800 apart
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Beryllium which act as the
center atom has the valence
electron of 2s2
b) Excited state
One of the electron from 2s ispromote to 2p orbitals - equal
energy level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of a
linear shape
Molecular shape Linear
Angle between bond pair 1800
4 The concept of hybridisation is also useful to explain molecules with
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p y p
doubletriple bonds By using the concept of the direct overlapping
orbitals and side-touch lapping orbitals the formation of multiple
bonds in ethene C2H4 and ethyne C2H2 are described
a) Ethene C2H4 - Hybridisation take place for both carbon atoms in
ethene molecule is sp2 hybridisation
bull When both hybridised C is bonded together one of the hybridisedorbital overlapped directly between each other while the other two
hybridised orbitals overlapped directly with hydrogen atoms
z
atoms form a side-touch bond between each other and form anotherbond as shown in the diagram below
bull From the diagram with C=C there are two types of bond A sigma-
bond (σ-bond) is covalent bonds formed by orbitals overlapping end-
to-end with the electron density concentrated between the nuclei ofthe bonding atoms while the second type is called a pi bond (π-
bond) which is defined as a covalent bond formed by sideways
overlapping orbitals with electron density concentrated above and
below the plane of the nuclei of the bonding atoms
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital of 2s and two
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of atrigonal planar shape
Note that on unhybridised pz
orbital an electron is presence
Molecular shape Trigonal planar Angle between bond pair 1200
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b) Ethyne C H has the Lewis structure of H CequivC H which the
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b) Ethyne C2H2 has the Lewis structure of HminusCequivCminusH which the
bonding can be explain using sp hybridisation Table below
described how the hybridisation take place on each carbonatom and how the formation of triple bond occur
c) From the diagram formation of -CequivC- is due to the formation
of one sigma-bond by direct overlapping of one of the twohybrid orbitals on each C while the other two pi-bond bonds
are formed as a result of side-lapping of the each two
y z
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital from 2s and 2p
orbitals combined (hybrid) and
rearrange themselves to the
shape and orientation of a
linear shapeNote that on unhybridised py amp
pz orbitals electrons are
presence to form two
π-bonds Molecular shape Linear Angle between bond pair 1800
sp
pz py
5 Other examples of applications in valence bond theory includes the
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formation of nitrogen molecule N2 and hydrogen cyanide HCN
molecule
a) Nitrogen gas is Earths most abundant gas as it cover 78 of the
content of our air Nitrogen molecule is an inert gas thanks to its
short covalent bond and also its strong triple bond Therefore
a lot of heats are required to break the chemical bond of nitrogenbefore it can be applied in industries Using valence bond theory
the bonding of nitrogen molecule is explained in the diagram
b) So when the two hybridised nitrogen atom interacting amongeach other they hence formed a linear shape and two pi-bonds
(πminusbonds) are formed as a result of side-lapping of unhybridised
py and pz orbital respectively
N N
ExplanationEnergy level
diagramDiagram of orbitals
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diagram
a) Ground state
Both nitrogen whichact as the center atom
has the valence
electron of 2s22p3
b) Excited state One of the electron
from 2s is promote to 2p
orbitals - e ual ener
level
c) Hybridised state
One orbital of 2s and
one orbitals of 2p
combined (hybrid) and
rearrange themselves
to the shape and
orientation of a linear
shapeMolecular shape Linear
Angle between bond pair 1800
b) Whereas for hydrogen cyanide HCN both carbon atom and
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) y g y
nitrogen atom undergoes sp hybridisation in order to form a
linear structureExplanation
Energy level diagram for
carbon atom C
Energy level diagram for
nitrogen atom N
a) Ground state
Valence electronC 2s22p2
N 2s22p3
b) Excited state
One of the electron from
2s is promote to 2p
orbitals - equal energy
level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined
(hybrid) and form sp
hybridisation
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N
σσσσ
π
σσσσ
π
6 However there are a few limitation on valence bond theory such
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as when explaining the effect angle of bond-pair and bond-pair
electrons when there isare presence of lone pair electron in thecenter atom and also the difference of electronegativity
a) Bonding in ammonia NH3 and water H2O ~ Both NH3 and
H2O undergoes an arrangement similar to sp3 hybridisation
similar to that of C in methane Table below shows thehybridisation occur for nitrogen in ammonia and oxygen in water
Energy level diagram Energy level diagram
for nitrogen atom N for oxygen atom O
a) Ground state
Valence electron
N 2s22p3
O 2s2
2p4
b) Excited state
One of the electron
from 2s is promote to
2p orbitals
c) Hybridised state
One orbital of 2s and
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One orbital of 2s and
three orbitals from 2p
combined (hybrid) andrearrange themselves to
the shape and
orientation similar to that
of tetrahedral shape
Similar to arrangement in
tetrahedral
Similar to arrangement in
tetrahedral
of ammonia NH3 andwater H2O
Shape and angleShape trigonal pyramidal
Angle 1070
Shape bent
Angle 10450
Number of bond pair amp
lone pair electrons
Bond pair electron 3
Lone pair electron 1
Bond pair electron 2
Lone pair electron 2
bull The angle of bond pair - bond pair electrons in ammonia and0 0
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water are 1070 and 10450 respectively which is lesser than in
methane molecule (10950
) This can be explained by the factof the presence of lone-pair electrons in both ammonia and
water Since the lone pair - lone pair electron repulsion is
stronger than lone pair - bond pair electron repulsion than
bond pair - bond pair electron repulsion it is expected that therepulsion between lone pair - bond pair electrons in ammonia
is stronger than bond pair - bond pair electrons repulsion
hence caused the angle to squeeze to a smaller angle As
for water since there is a presence of lone pair - lone pair
electrons repulsion it results the bond pair - bond pair
electron repulsion to be much smaller hence caused the
angle to squeeze to a smaller angle Though valence bondtheory does not actually explain the hybridisation especially
when it relates to the repulsion occur involving lone pair
electron another theories shall be applied to study such
effect All these shall be discussed further in VSEPR theory
b) Effect of electronegativity and bonding angle between bond pair
- bond pair electrons ~
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- bond pair electrons
i As discussed earlier in the bonding of water (H2
O) the
molecular shape and angle is described in the diagram below
Sul hur which is also an element from Grou 16 formed h dro en
sulphide H2S when sulphur react with hydrogen However unlike
water the bonding angles is much smaller compare to water This is
due to the difference of electronegativity and also the bond length
between O and S in molecule Since O is more electronegative
compare to S the O-H bond is pulled closer toward O Furthermorethe bond length of O-H is shorter compare to S-H As a result the
bonding pair-bonding pair electrons repulse greater with each other in
H-O-H and caused the angle become greater
ii Another example is between NH3 and PH3 The orbital
f
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diagram for both NH3 and PH3 are described below Both
nitrogen N and phosphorous P are from the same groupwhich is Group 15
Similar to the above case P in phosphine is less
electronegative than N in ammonia and bond length of N-H
is shorter than P-H As a result H is pull closer to N and
repulsion between bonding of H-N-H is greater compare toH-P-H and caused the angle between H-N-H is greater
compare to H-P-H
324 Valence Shell Electron Pair Repulsion (VSEPR) Theory
1 M l l t i th th di i l t f t
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1 Molecular geometry is the three-dimensional arrangement of atoms
in a molecule A moleculersquos geometry affects its physical andchemical properties such as melting point boiling point density
and the types of reactions it undergoes The basic concept of
VSEPR theory is based on the three general rules below
a) As far as electron-pair repulsion is concerned double bonds andtriple bonds can be treated like single bonds This approximation is
good for qualitative purposes However you should realize that in
ldquo rdquo
because there are two or three bonds between two atoms theelectron density occupies more space
b) If a molecule has two or more resonance structures we can apply
the VSEPR model to any one of them Formal charges are usually
not shown
c) The order of repulsion strength of lone pair and bond pair are
lone-pair amp lone-pair electrons repulsion are the strongest
followed by lone-pair amp bond-pair electrons repulsion while
bond-pair amp bond-pair electrons repulsion is the weakest
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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2 A Lewis structure is a representation of covalent bonding
in which shared electron pairs are shown either as lines
or as pairs of dots between two atoms and lone pairsare shown as pairs of dots on individual atoms Only
valence electrons are shown in a Lewis structure
a) Consider the fluorine molecule F2 The electronconfiguration of F is 1s22s22p5 The 1s electrons are in
inner shell which is nearest to the nucleus For this reason
the do not artici ate in bond formation Thus each F
atom has seven valence electrons (2s22p5) Thereforeeach fluorine atom needed one electron to achieve octet
configuration (ns2np6)
F 1 s22s22p5 F 1 s22s22p5 1s2 2s2 2p6 1s2 2s2 2p6
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bull Oxygen atom has electronic configuration of 1s22s22p4 To
achieve stable octet configuration (ns2np6) each oxygen
atom need 2 electrons Hence when 2 oxygen atomsinteract they shared two electrons in between each other as
described in diagram below
O 1s2
2s2
2p4
O 1s2
2s2
2p4
1s2
2s2
2p6
1s2
2s2
2p6
bull From the structure of oxygen molecule formed each oxygen
atom shared two electrons from each other to form a double
bond in order to achieve octet configuration among each
oxygen atom
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bull Nitrogen has electronic configuration of 1s22s22p3 and
required 3 electrons to achieve octet configuration (ns2np6)
In this case each nitrogen atom shared 3 electrons fromeach of its atom to form triple bond
N 1 s22s22p3 N 1 s22s22p3 1s2 2s2 2p6 1s2 2s2 2p6
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Water H2O
H 1 s1 O 2 s22p4
Carbon dioxide CO2
C 2 s22p2 O 2 s22p4
Ammonia NH3
H 1 s1 N 2 s22p3Ethene C2H4
H 1 s1 C 2 s22p2
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Hydrogen cyanide HCN
H 1 s1 N 2 s22p3 C 2 s22p2
Ethanoic acid CH3COOH
H 1 s1 O 2 s22p4 C 2 s22p2
Tetrachloromethane CCl4C 2 s22p2 Cl 3s23p5
Ethyne C2H2
H 1 s1 C 2 s22p2
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a) Note that in ethene hydrogen cyanide and ethanoic acid all
the valence electrons are used in bonding there are no lone
pairs on the carbon atoms In fact most of the stablemolecules containing carbon do not have lone pairs on the
carbon atoms
b) Multiple bonds (double bond or triple bond) are shorter than single covalent bonds Bond length is defined as the
distance between the nuclei of two covalently bonded atoms
nitrogen triple bonds are shorter than double bonds whichin turn are shorter than single bonds The shorter multiple
bonds are also more stable than single bonds
c) Covalent bond not only exist in neutral molecule but also insome molecular ions Table below shows a few example of
molecular ions which have covalent bonds in its molecule
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Carbonate ion CO32-
C 2 s22p2 O 2 s22p4
Cyanide ion CN-
N 2 s22p3 C 2 s22p2
u p a e on 4-
S 3 s23p4 O 2 s22p4
ra e on 3-
N 2 s22p3 O 2 s22p4
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321 Exception of Octet Rules
1 From the Lewis structure sketched for sulphate ion SO42- we can
see that the center metal atom (sulphur) has more than 8electrons For this the molecule is described as molecule that can
expanded octet These molecules that have more than 8
electrons are located at Period 3 and below as these center
atoms have empty d-orbital available to expand the number ofelectrons positioned in the center atom Those center atom from
Period 2 such as C N O and F can only allocate a maximum
Example Both phosphorous and nitrogen are elements from Group15 with the valence electron of ns2np3 They can both form NCl3and PCl3 respectively when react with limited amount of chlorine
however under excess chlorine only PCl 5 can be formed but
not NCl 5 Explain the statement bolded
Solution This is due to phosphorous which is from period
3 have empty d-orbital to expand the octet However
nitrogen which is from Period 2 do not have emptyorbital and can only allocate 8 electrons in its shell
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Phosphorous pentachloride PCl5P 3 s23p3 Cl 3s23p5
Sulphur hexafluoride SF6
S 3 s23p4 F 2 s22p5
Bromine pentachloride BrCl5Cl 3s23p5 Br 4s24p5
Xenon tetrafluoride XeF4
Xe 5s25p6 F 2 s22p5
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2 There are also some stable covalent compounds which have
less than 8 electrons in their center atom (incomplete octet)
The center atoms are usually metals with great number ofvalence electrons with small atomic radius such as
beryllium boron and aluminium
BeCl2Be 2s2 Cl 3s23p5 BF3B 2 s22p1 F 2 s22p5 AlCl3Al 3s23p1 F 3 s23p5
These compound possessed stability due to their short bondlength between center atom with surround atom Furthermore
they can form resonance structure between the center atom
and surrounding atoms
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3 There are also some molecules which contain an odd
number of electrons Among the most common compounds
are nitrogen monoxide (NO) and nitrogen dioxide (NO2)
Odd-electron molecules are sometimes called radicals
Nitrogen monoxide NO
N 2 s22p3 O 2 s22p4
Nitrogen dioxide NO2
N 2 s22p3 O 2 s22p4
any ra ca s are g y reac ve e reason s t at t ere
is a tendency for the unpaired electron to form a
covalent bond with an unpaired electron on another
molecule For example when two nitrogen dioxide
molecules collide they form dinitrogen tetroxide
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322 Dative bond
1 As shown in the Lewis structure of nitrate ion one of the N-Obond is drawn as rarr The bond rarr placed is called as dative
bond (also known as coordinative bond) where dative
bond is defined as a covalent bond in which one of the atoms
donate the lone pair electrons available Although theproperties of a coordinate covalent bond do not differ from
those of a normal covalent bond because all electrons are
a e no ma er w a e r source
2 Dative bond is usually applied for these few circumstances
below
a) To assist atom molecule ion that not yet achieved octet
configuration Making use of atom which has lone pair
electrons to those which are lack of electrons
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Sulphur dioxide SO2Sulphur trioxide SO3 Carbon monoxide CO
Ozone molecule O3 Water with hydrogen ion
S
OO
O
Ammonia with hydrogen ion Ammonia with boron trifluoride
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b) Formation of dimer - In order for some compounds which
have incomplete octet to achieve stability they tend to form
dimer or polymer among themselves by using dative bondTwo of the most common examples are aluminium trichloride
and beryllium dichloride
Monomer of AlCl3Dimer of aluminium chloride Al
2
Cl6
Monomer of BeCl2 Dimer of Be2Cl4 Polymer of (BeCl2)n
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c) Formation of coordination compounds - Coordination
compounds are substances that contain at least one complex
ion a species consisting of a central metal cation (either atransition metal or a main-group metal) that is bonded to
molecules andor anions called ligands via dative
(coordinative) bond
hexaaquacopper (II) ion [Cu(H2O)6]2+ tetraamminenickel (II) ion [Ni(NH3)4]
2+
Hexacyanoferrate (III) ion [Fe(CN)6]3- Trioxalatocobaltate(III)ion [Co(C2O4)3]
3-
3 2 3 Hybridisation Theory
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323 Hybridisation Theory
Valence bond theory
bull The basic principle of valence bond theory is that a covalentbond is formed when orbitals of two atoms overlap and the
overlapped region which is between the nuclei is occupied by
a pair of electrons The central themes of valence bond theory
derive from this principle
ndash Opposing spins of the electron pair ~ Stated in Paulis
orbitals has a maximum capacity of two electrons that musthave opposite spins For example when a covalent bond is
formed in molecule of hydrogen H2 the two 1s electrons of
two H atoms occupy the overlapping 1s orbitals and have
opposite spins
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bull Maximum overlap of bonding orbitals ~ The bond strength
depends on attraction of the nuclei for the shared electrons
so the greater the orbital overlap the stronger (morestable) the bond The extent of overlap depends on the
shapes and directions of the orbitals An s orbital is spherical
but p and d orbitals have more electron density in one
direction than in another Thus whenever possible a bond
involving p or d orbitals will be oriented in the direction that
maximizes overlap For example in hydrogen fluoride (HF)
bond the 1s orbital of H overlaps the half-filled 2p orbital of Falong the long axis of that orbital Any other direction would
result in less overlap and thus a weaker bond
H b idi ti f t i bit l T t f th
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bull Hybridisation of atomic orbitals ~ To account for the
bonding in simple diatomic molecules like HF we picture the
direct overlap of s and p orbitals of isolated atoms But howcan we account for the shapes of so many molecules and
polyatomic ions through the overlap of spherical s orbitals
dumbbell-shaped p orbitals and cloverleaf-shaped d orbitals
Linus Pauling proposed that the valence atomic orbitals in the
molecule are different from those in the isolated atoms The
spatial orientations of these new orbitals lead to more stable
bonds and are consistent with observed molecular shapesThe process of orbital mixing is called hybridisation and
the new atomic orbitals are called hybrid orbitals Two key
points about the number and type of hybrid orbitals are that
i The number of hybrid orbitals obtained equals the number ofatomic orbitals mixed
ii The type of hybrid orbitals obtained varies with the types of
atomic orbitals mixed
3 2 3 1 T f h b idi ti
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3231 Type of hybridisation
1 sp3 hybridisation ~ When four electron groups surround
the central atom the center atom involved must prepare fourorbitals with equal energies to overlap with the four
surrounding electron groups Valence Bond theory uses
hypothetical hybrid orbitals which are atomic orbitals
obtained when two or more non-equivalent orbitals of the
same atom combine in preparation for covalent bond
formation H bridisation is the term a lied to the mixin of
atomic orbitals in an atom (usually a central atom) togenerate a set of hybrid orbitals In the case of sp3 We can
generate four equivalent hybrid orbitals from the center
atom by mixing the s orbital and the three p orbitals
Explanation Energy level diagram Diagram of orbitals
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Explanation Energy level diagram Diagram of orbitals
a) Ground state
Carbon which act as thecenter atom has the valence
electron of 2s22p2
b) Excited state One of the electron from 2s
is promote to 2p orbitals -
equa energy eve
c) Hybridised state
One orbital of 2s and three
orbitals of 2p combined(hybrid) and rearrange
themselves to the shape
and orientation of a
tetrahedral shape
Molecular shape Tetrahedral
Angle between bond pair 10950
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2 sp2 hybridisation ~ When three electron groups surround
the central atom the center atom involved must prepare three
orbitals with equal energies to overlap with the three
surrounding electron groups In sp2 hybridisation three
equivalent (in terms of energy level) from the center atom by
mixing the one s orbital and the two p orbitals Using boron
trifluoride (BF3) as example sp2
hybridisation is explained
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3 sp hybridisation ~ When two electron groups surround
the central atom we observe a linear shape which
means that the bonding orbitals must have a linear
orientation VB theory explains this by proposing thatmixing two nonequivalent orbitals of a central atom one
s and one p gives rise to two equivalent sp hybrid
orbitals that lie 1800 apart
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Beryllium which act as the
center atom has the valence
electron of 2s2
b) Excited state
One of the electron from 2s ispromote to 2p orbitals - equal
energy level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of a
linear shape
Molecular shape Linear
Angle between bond pair 1800
4 The concept of hybridisation is also useful to explain molecules with
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p y p
doubletriple bonds By using the concept of the direct overlapping
orbitals and side-touch lapping orbitals the formation of multiple
bonds in ethene C2H4 and ethyne C2H2 are described
a) Ethene C2H4 - Hybridisation take place for both carbon atoms in
ethene molecule is sp2 hybridisation
bull When both hybridised C is bonded together one of the hybridisedorbital overlapped directly between each other while the other two
hybridised orbitals overlapped directly with hydrogen atoms
z
atoms form a side-touch bond between each other and form anotherbond as shown in the diagram below
bull From the diagram with C=C there are two types of bond A sigma-
bond (σ-bond) is covalent bonds formed by orbitals overlapping end-
to-end with the electron density concentrated between the nuclei ofthe bonding atoms while the second type is called a pi bond (π-
bond) which is defined as a covalent bond formed by sideways
overlapping orbitals with electron density concentrated above and
below the plane of the nuclei of the bonding atoms
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital of 2s and two
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of atrigonal planar shape
Note that on unhybridised pz
orbital an electron is presence
Molecular shape Trigonal planar Angle between bond pair 1200
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b) Ethyne C H has the Lewis structure of H CequivC H which the
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b) Ethyne C2H2 has the Lewis structure of HminusCequivCminusH which the
bonding can be explain using sp hybridisation Table below
described how the hybridisation take place on each carbonatom and how the formation of triple bond occur
c) From the diagram formation of -CequivC- is due to the formation
of one sigma-bond by direct overlapping of one of the twohybrid orbitals on each C while the other two pi-bond bonds
are formed as a result of side-lapping of the each two
y z
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital from 2s and 2p
orbitals combined (hybrid) and
rearrange themselves to the
shape and orientation of a
linear shapeNote that on unhybridised py amp
pz orbitals electrons are
presence to form two
π-bonds Molecular shape Linear Angle between bond pair 1800
sp
pz py
5 Other examples of applications in valence bond theory includes the
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formation of nitrogen molecule N2 and hydrogen cyanide HCN
molecule
a) Nitrogen gas is Earths most abundant gas as it cover 78 of the
content of our air Nitrogen molecule is an inert gas thanks to its
short covalent bond and also its strong triple bond Therefore
a lot of heats are required to break the chemical bond of nitrogenbefore it can be applied in industries Using valence bond theory
the bonding of nitrogen molecule is explained in the diagram
b) So when the two hybridised nitrogen atom interacting amongeach other they hence formed a linear shape and two pi-bonds
(πminusbonds) are formed as a result of side-lapping of unhybridised
py and pz orbital respectively
N N
ExplanationEnergy level
diagramDiagram of orbitals
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diagram
a) Ground state
Both nitrogen whichact as the center atom
has the valence
electron of 2s22p3
b) Excited state One of the electron
from 2s is promote to 2p
orbitals - e ual ener
level
c) Hybridised state
One orbital of 2s and
one orbitals of 2p
combined (hybrid) and
rearrange themselves
to the shape and
orientation of a linear
shapeMolecular shape Linear
Angle between bond pair 1800
b) Whereas for hydrogen cyanide HCN both carbon atom and
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) y g y
nitrogen atom undergoes sp hybridisation in order to form a
linear structureExplanation
Energy level diagram for
carbon atom C
Energy level diagram for
nitrogen atom N
a) Ground state
Valence electronC 2s22p2
N 2s22p3
b) Excited state
One of the electron from
2s is promote to 2p
orbitals - equal energy
level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined
(hybrid) and form sp
hybridisation
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N
σσσσ
π
σσσσ
π
6 However there are a few limitation on valence bond theory such
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as when explaining the effect angle of bond-pair and bond-pair
electrons when there isare presence of lone pair electron in thecenter atom and also the difference of electronegativity
a) Bonding in ammonia NH3 and water H2O ~ Both NH3 and
H2O undergoes an arrangement similar to sp3 hybridisation
similar to that of C in methane Table below shows thehybridisation occur for nitrogen in ammonia and oxygen in water
Energy level diagram Energy level diagram
for nitrogen atom N for oxygen atom O
a) Ground state
Valence electron
N 2s22p3
O 2s2
2p4
b) Excited state
One of the electron
from 2s is promote to
2p orbitals
c) Hybridised state
One orbital of 2s and
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One orbital of 2s and
three orbitals from 2p
combined (hybrid) andrearrange themselves to
the shape and
orientation similar to that
of tetrahedral shape
Similar to arrangement in
tetrahedral
Similar to arrangement in
tetrahedral
of ammonia NH3 andwater H2O
Shape and angleShape trigonal pyramidal
Angle 1070
Shape bent
Angle 10450
Number of bond pair amp
lone pair electrons
Bond pair electron 3
Lone pair electron 1
Bond pair electron 2
Lone pair electron 2
bull The angle of bond pair - bond pair electrons in ammonia and0 0
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water are 1070 and 10450 respectively which is lesser than in
methane molecule (10950
) This can be explained by the factof the presence of lone-pair electrons in both ammonia and
water Since the lone pair - lone pair electron repulsion is
stronger than lone pair - bond pair electron repulsion than
bond pair - bond pair electron repulsion it is expected that therepulsion between lone pair - bond pair electrons in ammonia
is stronger than bond pair - bond pair electrons repulsion
hence caused the angle to squeeze to a smaller angle As
for water since there is a presence of lone pair - lone pair
electrons repulsion it results the bond pair - bond pair
electron repulsion to be much smaller hence caused the
angle to squeeze to a smaller angle Though valence bondtheory does not actually explain the hybridisation especially
when it relates to the repulsion occur involving lone pair
electron another theories shall be applied to study such
effect All these shall be discussed further in VSEPR theory
b) Effect of electronegativity and bonding angle between bond pair
- bond pair electrons ~
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- bond pair electrons
i As discussed earlier in the bonding of water (H2
O) the
molecular shape and angle is described in the diagram below
Sul hur which is also an element from Grou 16 formed h dro en
sulphide H2S when sulphur react with hydrogen However unlike
water the bonding angles is much smaller compare to water This is
due to the difference of electronegativity and also the bond length
between O and S in molecule Since O is more electronegative
compare to S the O-H bond is pulled closer toward O Furthermorethe bond length of O-H is shorter compare to S-H As a result the
bonding pair-bonding pair electrons repulse greater with each other in
H-O-H and caused the angle become greater
ii Another example is between NH3 and PH3 The orbital
f
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diagram for both NH3 and PH3 are described below Both
nitrogen N and phosphorous P are from the same groupwhich is Group 15
Similar to the above case P in phosphine is less
electronegative than N in ammonia and bond length of N-H
is shorter than P-H As a result H is pull closer to N and
repulsion between bonding of H-N-H is greater compare toH-P-H and caused the angle between H-N-H is greater
compare to H-P-H
324 Valence Shell Electron Pair Repulsion (VSEPR) Theory
1 M l l t i th th di i l t f t
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1 Molecular geometry is the three-dimensional arrangement of atoms
in a molecule A moleculersquos geometry affects its physical andchemical properties such as melting point boiling point density
and the types of reactions it undergoes The basic concept of
VSEPR theory is based on the three general rules below
a) As far as electron-pair repulsion is concerned double bonds andtriple bonds can be treated like single bonds This approximation is
good for qualitative purposes However you should realize that in
ldquo rdquo
because there are two or three bonds between two atoms theelectron density occupies more space
b) If a molecule has two or more resonance structures we can apply
the VSEPR model to any one of them Formal charges are usually
not shown
c) The order of repulsion strength of lone pair and bond pair are
lone-pair amp lone-pair electrons repulsion are the strongest
followed by lone-pair amp bond-pair electrons repulsion while
bond-pair amp bond-pair electrons repulsion is the weakest
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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bull Oxygen atom has electronic configuration of 1s22s22p4 To
achieve stable octet configuration (ns2np6) each oxygen
atom need 2 electrons Hence when 2 oxygen atomsinteract they shared two electrons in between each other as
described in diagram below
O 1s2
2s2
2p4
O 1s2
2s2
2p4
1s2
2s2
2p6
1s2
2s2
2p6
bull From the structure of oxygen molecule formed each oxygen
atom shared two electrons from each other to form a double
bond in order to achieve octet configuration among each
oxygen atom
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bull Nitrogen has electronic configuration of 1s22s22p3 and
required 3 electrons to achieve octet configuration (ns2np6)
In this case each nitrogen atom shared 3 electrons fromeach of its atom to form triple bond
N 1 s22s22p3 N 1 s22s22p3 1s2 2s2 2p6 1s2 2s2 2p6
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Water H2O
H 1 s1 O 2 s22p4
Carbon dioxide CO2
C 2 s22p2 O 2 s22p4
Ammonia NH3
H 1 s1 N 2 s22p3Ethene C2H4
H 1 s1 C 2 s22p2
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Hydrogen cyanide HCN
H 1 s1 N 2 s22p3 C 2 s22p2
Ethanoic acid CH3COOH
H 1 s1 O 2 s22p4 C 2 s22p2
Tetrachloromethane CCl4C 2 s22p2 Cl 3s23p5
Ethyne C2H2
H 1 s1 C 2 s22p2
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a) Note that in ethene hydrogen cyanide and ethanoic acid all
the valence electrons are used in bonding there are no lone
pairs on the carbon atoms In fact most of the stablemolecules containing carbon do not have lone pairs on the
carbon atoms
b) Multiple bonds (double bond or triple bond) are shorter than single covalent bonds Bond length is defined as the
distance between the nuclei of two covalently bonded atoms
nitrogen triple bonds are shorter than double bonds whichin turn are shorter than single bonds The shorter multiple
bonds are also more stable than single bonds
c) Covalent bond not only exist in neutral molecule but also insome molecular ions Table below shows a few example of
molecular ions which have covalent bonds in its molecule
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Carbonate ion CO32-
C 2 s22p2 O 2 s22p4
Cyanide ion CN-
N 2 s22p3 C 2 s22p2
u p a e on 4-
S 3 s23p4 O 2 s22p4
ra e on 3-
N 2 s22p3 O 2 s22p4
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321 Exception of Octet Rules
1 From the Lewis structure sketched for sulphate ion SO42- we can
see that the center metal atom (sulphur) has more than 8electrons For this the molecule is described as molecule that can
expanded octet These molecules that have more than 8
electrons are located at Period 3 and below as these center
atoms have empty d-orbital available to expand the number ofelectrons positioned in the center atom Those center atom from
Period 2 such as C N O and F can only allocate a maximum
Example Both phosphorous and nitrogen are elements from Group15 with the valence electron of ns2np3 They can both form NCl3and PCl3 respectively when react with limited amount of chlorine
however under excess chlorine only PCl 5 can be formed but
not NCl 5 Explain the statement bolded
Solution This is due to phosphorous which is from period
3 have empty d-orbital to expand the octet However
nitrogen which is from Period 2 do not have emptyorbital and can only allocate 8 electrons in its shell
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Phosphorous pentachloride PCl5P 3 s23p3 Cl 3s23p5
Sulphur hexafluoride SF6
S 3 s23p4 F 2 s22p5
Bromine pentachloride BrCl5Cl 3s23p5 Br 4s24p5
Xenon tetrafluoride XeF4
Xe 5s25p6 F 2 s22p5
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2 There are also some stable covalent compounds which have
less than 8 electrons in their center atom (incomplete octet)
The center atoms are usually metals with great number ofvalence electrons with small atomic radius such as
beryllium boron and aluminium
BeCl2Be 2s2 Cl 3s23p5 BF3B 2 s22p1 F 2 s22p5 AlCl3Al 3s23p1 F 3 s23p5
These compound possessed stability due to their short bondlength between center atom with surround atom Furthermore
they can form resonance structure between the center atom
and surrounding atoms
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3 There are also some molecules which contain an odd
number of electrons Among the most common compounds
are nitrogen monoxide (NO) and nitrogen dioxide (NO2)
Odd-electron molecules are sometimes called radicals
Nitrogen monoxide NO
N 2 s22p3 O 2 s22p4
Nitrogen dioxide NO2
N 2 s22p3 O 2 s22p4
any ra ca s are g y reac ve e reason s t at t ere
is a tendency for the unpaired electron to form a
covalent bond with an unpaired electron on another
molecule For example when two nitrogen dioxide
molecules collide they form dinitrogen tetroxide
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322 Dative bond
1 As shown in the Lewis structure of nitrate ion one of the N-Obond is drawn as rarr The bond rarr placed is called as dative
bond (also known as coordinative bond) where dative
bond is defined as a covalent bond in which one of the atoms
donate the lone pair electrons available Although theproperties of a coordinate covalent bond do not differ from
those of a normal covalent bond because all electrons are
a e no ma er w a e r source
2 Dative bond is usually applied for these few circumstances
below
a) To assist atom molecule ion that not yet achieved octet
configuration Making use of atom which has lone pair
electrons to those which are lack of electrons
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Sulphur dioxide SO2Sulphur trioxide SO3 Carbon monoxide CO
Ozone molecule O3 Water with hydrogen ion
S
OO
O
Ammonia with hydrogen ion Ammonia with boron trifluoride
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b) Formation of dimer - In order for some compounds which
have incomplete octet to achieve stability they tend to form
dimer or polymer among themselves by using dative bondTwo of the most common examples are aluminium trichloride
and beryllium dichloride
Monomer of AlCl3Dimer of aluminium chloride Al
2
Cl6
Monomer of BeCl2 Dimer of Be2Cl4 Polymer of (BeCl2)n
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c) Formation of coordination compounds - Coordination
compounds are substances that contain at least one complex
ion a species consisting of a central metal cation (either atransition metal or a main-group metal) that is bonded to
molecules andor anions called ligands via dative
(coordinative) bond
hexaaquacopper (II) ion [Cu(H2O)6]2+ tetraamminenickel (II) ion [Ni(NH3)4]
2+
Hexacyanoferrate (III) ion [Fe(CN)6]3- Trioxalatocobaltate(III)ion [Co(C2O4)3]
3-
3 2 3 Hybridisation Theory
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323 Hybridisation Theory
Valence bond theory
bull The basic principle of valence bond theory is that a covalentbond is formed when orbitals of two atoms overlap and the
overlapped region which is between the nuclei is occupied by
a pair of electrons The central themes of valence bond theory
derive from this principle
ndash Opposing spins of the electron pair ~ Stated in Paulis
orbitals has a maximum capacity of two electrons that musthave opposite spins For example when a covalent bond is
formed in molecule of hydrogen H2 the two 1s electrons of
two H atoms occupy the overlapping 1s orbitals and have
opposite spins
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bull Maximum overlap of bonding orbitals ~ The bond strength
depends on attraction of the nuclei for the shared electrons
so the greater the orbital overlap the stronger (morestable) the bond The extent of overlap depends on the
shapes and directions of the orbitals An s orbital is spherical
but p and d orbitals have more electron density in one
direction than in another Thus whenever possible a bond
involving p or d orbitals will be oriented in the direction that
maximizes overlap For example in hydrogen fluoride (HF)
bond the 1s orbital of H overlaps the half-filled 2p orbital of Falong the long axis of that orbital Any other direction would
result in less overlap and thus a weaker bond
H b idi ti f t i bit l T t f th
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bull Hybridisation of atomic orbitals ~ To account for the
bonding in simple diatomic molecules like HF we picture the
direct overlap of s and p orbitals of isolated atoms But howcan we account for the shapes of so many molecules and
polyatomic ions through the overlap of spherical s orbitals
dumbbell-shaped p orbitals and cloverleaf-shaped d orbitals
Linus Pauling proposed that the valence atomic orbitals in the
molecule are different from those in the isolated atoms The
spatial orientations of these new orbitals lead to more stable
bonds and are consistent with observed molecular shapesThe process of orbital mixing is called hybridisation and
the new atomic orbitals are called hybrid orbitals Two key
points about the number and type of hybrid orbitals are that
i The number of hybrid orbitals obtained equals the number ofatomic orbitals mixed
ii The type of hybrid orbitals obtained varies with the types of
atomic orbitals mixed
3 2 3 1 T f h b idi ti
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3231 Type of hybridisation
1 sp3 hybridisation ~ When four electron groups surround
the central atom the center atom involved must prepare fourorbitals with equal energies to overlap with the four
surrounding electron groups Valence Bond theory uses
hypothetical hybrid orbitals which are atomic orbitals
obtained when two or more non-equivalent orbitals of the
same atom combine in preparation for covalent bond
formation H bridisation is the term a lied to the mixin of
atomic orbitals in an atom (usually a central atom) togenerate a set of hybrid orbitals In the case of sp3 We can
generate four equivalent hybrid orbitals from the center
atom by mixing the s orbital and the three p orbitals
Explanation Energy level diagram Diagram of orbitals
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Explanation Energy level diagram Diagram of orbitals
a) Ground state
Carbon which act as thecenter atom has the valence
electron of 2s22p2
b) Excited state One of the electron from 2s
is promote to 2p orbitals -
equa energy eve
c) Hybridised state
One orbital of 2s and three
orbitals of 2p combined(hybrid) and rearrange
themselves to the shape
and orientation of a
tetrahedral shape
Molecular shape Tetrahedral
Angle between bond pair 10950
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2 sp2 hybridisation ~ When three electron groups surround
the central atom the center atom involved must prepare three
orbitals with equal energies to overlap with the three
surrounding electron groups In sp2 hybridisation three
equivalent (in terms of energy level) from the center atom by
mixing the one s orbital and the two p orbitals Using boron
trifluoride (BF3) as example sp2
hybridisation is explained
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3 sp hybridisation ~ When two electron groups surround
the central atom we observe a linear shape which
means that the bonding orbitals must have a linear
orientation VB theory explains this by proposing thatmixing two nonequivalent orbitals of a central atom one
s and one p gives rise to two equivalent sp hybrid
orbitals that lie 1800 apart
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Beryllium which act as the
center atom has the valence
electron of 2s2
b) Excited state
One of the electron from 2s ispromote to 2p orbitals - equal
energy level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of a
linear shape
Molecular shape Linear
Angle between bond pair 1800
4 The concept of hybridisation is also useful to explain molecules with
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p y p
doubletriple bonds By using the concept of the direct overlapping
orbitals and side-touch lapping orbitals the formation of multiple
bonds in ethene C2H4 and ethyne C2H2 are described
a) Ethene C2H4 - Hybridisation take place for both carbon atoms in
ethene molecule is sp2 hybridisation
bull When both hybridised C is bonded together one of the hybridisedorbital overlapped directly between each other while the other two
hybridised orbitals overlapped directly with hydrogen atoms
z
atoms form a side-touch bond between each other and form anotherbond as shown in the diagram below
bull From the diagram with C=C there are two types of bond A sigma-
bond (σ-bond) is covalent bonds formed by orbitals overlapping end-
to-end with the electron density concentrated between the nuclei ofthe bonding atoms while the second type is called a pi bond (π-
bond) which is defined as a covalent bond formed by sideways
overlapping orbitals with electron density concentrated above and
below the plane of the nuclei of the bonding atoms
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital of 2s and two
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of atrigonal planar shape
Note that on unhybridised pz
orbital an electron is presence
Molecular shape Trigonal planar Angle between bond pair 1200
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b) Ethyne C H has the Lewis structure of H CequivC H which the
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b) Ethyne C2H2 has the Lewis structure of HminusCequivCminusH which the
bonding can be explain using sp hybridisation Table below
described how the hybridisation take place on each carbonatom and how the formation of triple bond occur
c) From the diagram formation of -CequivC- is due to the formation
of one sigma-bond by direct overlapping of one of the twohybrid orbitals on each C while the other two pi-bond bonds
are formed as a result of side-lapping of the each two
y z
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital from 2s and 2p
orbitals combined (hybrid) and
rearrange themselves to the
shape and orientation of a
linear shapeNote that on unhybridised py amp
pz orbitals electrons are
presence to form two
π-bonds Molecular shape Linear Angle between bond pair 1800
sp
pz py
5 Other examples of applications in valence bond theory includes the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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formation of nitrogen molecule N2 and hydrogen cyanide HCN
molecule
a) Nitrogen gas is Earths most abundant gas as it cover 78 of the
content of our air Nitrogen molecule is an inert gas thanks to its
short covalent bond and also its strong triple bond Therefore
a lot of heats are required to break the chemical bond of nitrogenbefore it can be applied in industries Using valence bond theory
the bonding of nitrogen molecule is explained in the diagram
b) So when the two hybridised nitrogen atom interacting amongeach other they hence formed a linear shape and two pi-bonds
(πminusbonds) are formed as a result of side-lapping of unhybridised
py and pz orbital respectively
N N
ExplanationEnergy level
diagramDiagram of orbitals
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diagram
a) Ground state
Both nitrogen whichact as the center atom
has the valence
electron of 2s22p3
b) Excited state One of the electron
from 2s is promote to 2p
orbitals - e ual ener
level
c) Hybridised state
One orbital of 2s and
one orbitals of 2p
combined (hybrid) and
rearrange themselves
to the shape and
orientation of a linear
shapeMolecular shape Linear
Angle between bond pair 1800
b) Whereas for hydrogen cyanide HCN both carbon atom and
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) y g y
nitrogen atom undergoes sp hybridisation in order to form a
linear structureExplanation
Energy level diagram for
carbon atom C
Energy level diagram for
nitrogen atom N
a) Ground state
Valence electronC 2s22p2
N 2s22p3
b) Excited state
One of the electron from
2s is promote to 2p
orbitals - equal energy
level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined
(hybrid) and form sp
hybridisation
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N
σσσσ
π
σσσσ
π
6 However there are a few limitation on valence bond theory such
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as when explaining the effect angle of bond-pair and bond-pair
electrons when there isare presence of lone pair electron in thecenter atom and also the difference of electronegativity
a) Bonding in ammonia NH3 and water H2O ~ Both NH3 and
H2O undergoes an arrangement similar to sp3 hybridisation
similar to that of C in methane Table below shows thehybridisation occur for nitrogen in ammonia and oxygen in water
Energy level diagram Energy level diagram
for nitrogen atom N for oxygen atom O
a) Ground state
Valence electron
N 2s22p3
O 2s2
2p4
b) Excited state
One of the electron
from 2s is promote to
2p orbitals
c) Hybridised state
One orbital of 2s and
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One orbital of 2s and
three orbitals from 2p
combined (hybrid) andrearrange themselves to
the shape and
orientation similar to that
of tetrahedral shape
Similar to arrangement in
tetrahedral
Similar to arrangement in
tetrahedral
of ammonia NH3 andwater H2O
Shape and angleShape trigonal pyramidal
Angle 1070
Shape bent
Angle 10450
Number of bond pair amp
lone pair electrons
Bond pair electron 3
Lone pair electron 1
Bond pair electron 2
Lone pair electron 2
bull The angle of bond pair - bond pair electrons in ammonia and0 0
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water are 1070 and 10450 respectively which is lesser than in
methane molecule (10950
) This can be explained by the factof the presence of lone-pair electrons in both ammonia and
water Since the lone pair - lone pair electron repulsion is
stronger than lone pair - bond pair electron repulsion than
bond pair - bond pair electron repulsion it is expected that therepulsion between lone pair - bond pair electrons in ammonia
is stronger than bond pair - bond pair electrons repulsion
hence caused the angle to squeeze to a smaller angle As
for water since there is a presence of lone pair - lone pair
electrons repulsion it results the bond pair - bond pair
electron repulsion to be much smaller hence caused the
angle to squeeze to a smaller angle Though valence bondtheory does not actually explain the hybridisation especially
when it relates to the repulsion occur involving lone pair
electron another theories shall be applied to study such
effect All these shall be discussed further in VSEPR theory
b) Effect of electronegativity and bonding angle between bond pair
- bond pair electrons ~
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- bond pair electrons
i As discussed earlier in the bonding of water (H2
O) the
molecular shape and angle is described in the diagram below
Sul hur which is also an element from Grou 16 formed h dro en
sulphide H2S when sulphur react with hydrogen However unlike
water the bonding angles is much smaller compare to water This is
due to the difference of electronegativity and also the bond length
between O and S in molecule Since O is more electronegative
compare to S the O-H bond is pulled closer toward O Furthermorethe bond length of O-H is shorter compare to S-H As a result the
bonding pair-bonding pair electrons repulse greater with each other in
H-O-H and caused the angle become greater
ii Another example is between NH3 and PH3 The orbital
f
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diagram for both NH3 and PH3 are described below Both
nitrogen N and phosphorous P are from the same groupwhich is Group 15
Similar to the above case P in phosphine is less
electronegative than N in ammonia and bond length of N-H
is shorter than P-H As a result H is pull closer to N and
repulsion between bonding of H-N-H is greater compare toH-P-H and caused the angle between H-N-H is greater
compare to H-P-H
324 Valence Shell Electron Pair Repulsion (VSEPR) Theory
1 M l l t i th th di i l t f t
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1 Molecular geometry is the three-dimensional arrangement of atoms
in a molecule A moleculersquos geometry affects its physical andchemical properties such as melting point boiling point density
and the types of reactions it undergoes The basic concept of
VSEPR theory is based on the three general rules below
a) As far as electron-pair repulsion is concerned double bonds andtriple bonds can be treated like single bonds This approximation is
good for qualitative purposes However you should realize that in
ldquo rdquo
because there are two or three bonds between two atoms theelectron density occupies more space
b) If a molecule has two or more resonance structures we can apply
the VSEPR model to any one of them Formal charges are usually
not shown
c) The order of repulsion strength of lone pair and bond pair are
lone-pair amp lone-pair electrons repulsion are the strongest
followed by lone-pair amp bond-pair electrons repulsion while
bond-pair amp bond-pair electrons repulsion is the weakest
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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bull Nitrogen has electronic configuration of 1s22s22p3 and
required 3 electrons to achieve octet configuration (ns2np6)
In this case each nitrogen atom shared 3 electrons fromeach of its atom to form triple bond
N 1 s22s22p3 N 1 s22s22p3 1s2 2s2 2p6 1s2 2s2 2p6
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Water H2O
H 1 s1 O 2 s22p4
Carbon dioxide CO2
C 2 s22p2 O 2 s22p4
Ammonia NH3
H 1 s1 N 2 s22p3Ethene C2H4
H 1 s1 C 2 s22p2
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Hydrogen cyanide HCN
H 1 s1 N 2 s22p3 C 2 s22p2
Ethanoic acid CH3COOH
H 1 s1 O 2 s22p4 C 2 s22p2
Tetrachloromethane CCl4C 2 s22p2 Cl 3s23p5
Ethyne C2H2
H 1 s1 C 2 s22p2
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a) Note that in ethene hydrogen cyanide and ethanoic acid all
the valence electrons are used in bonding there are no lone
pairs on the carbon atoms In fact most of the stablemolecules containing carbon do not have lone pairs on the
carbon atoms
b) Multiple bonds (double bond or triple bond) are shorter than single covalent bonds Bond length is defined as the
distance between the nuclei of two covalently bonded atoms
nitrogen triple bonds are shorter than double bonds whichin turn are shorter than single bonds The shorter multiple
bonds are also more stable than single bonds
c) Covalent bond not only exist in neutral molecule but also insome molecular ions Table below shows a few example of
molecular ions which have covalent bonds in its molecule
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Carbonate ion CO32-
C 2 s22p2 O 2 s22p4
Cyanide ion CN-
N 2 s22p3 C 2 s22p2
u p a e on 4-
S 3 s23p4 O 2 s22p4
ra e on 3-
N 2 s22p3 O 2 s22p4
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321 Exception of Octet Rules
1 From the Lewis structure sketched for sulphate ion SO42- we can
see that the center metal atom (sulphur) has more than 8electrons For this the molecule is described as molecule that can
expanded octet These molecules that have more than 8
electrons are located at Period 3 and below as these center
atoms have empty d-orbital available to expand the number ofelectrons positioned in the center atom Those center atom from
Period 2 such as C N O and F can only allocate a maximum
Example Both phosphorous and nitrogen are elements from Group15 with the valence electron of ns2np3 They can both form NCl3and PCl3 respectively when react with limited amount of chlorine
however under excess chlorine only PCl 5 can be formed but
not NCl 5 Explain the statement bolded
Solution This is due to phosphorous which is from period
3 have empty d-orbital to expand the octet However
nitrogen which is from Period 2 do not have emptyorbital and can only allocate 8 electrons in its shell
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Phosphorous pentachloride PCl5P 3 s23p3 Cl 3s23p5
Sulphur hexafluoride SF6
S 3 s23p4 F 2 s22p5
Bromine pentachloride BrCl5Cl 3s23p5 Br 4s24p5
Xenon tetrafluoride XeF4
Xe 5s25p6 F 2 s22p5
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2 There are also some stable covalent compounds which have
less than 8 electrons in their center atom (incomplete octet)
The center atoms are usually metals with great number ofvalence electrons with small atomic radius such as
beryllium boron and aluminium
BeCl2Be 2s2 Cl 3s23p5 BF3B 2 s22p1 F 2 s22p5 AlCl3Al 3s23p1 F 3 s23p5
These compound possessed stability due to their short bondlength between center atom with surround atom Furthermore
they can form resonance structure between the center atom
and surrounding atoms
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3 There are also some molecules which contain an odd
number of electrons Among the most common compounds
are nitrogen monoxide (NO) and nitrogen dioxide (NO2)
Odd-electron molecules are sometimes called radicals
Nitrogen monoxide NO
N 2 s22p3 O 2 s22p4
Nitrogen dioxide NO2
N 2 s22p3 O 2 s22p4
any ra ca s are g y reac ve e reason s t at t ere
is a tendency for the unpaired electron to form a
covalent bond with an unpaired electron on another
molecule For example when two nitrogen dioxide
molecules collide they form dinitrogen tetroxide
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322 Dative bond
1 As shown in the Lewis structure of nitrate ion one of the N-Obond is drawn as rarr The bond rarr placed is called as dative
bond (also known as coordinative bond) where dative
bond is defined as a covalent bond in which one of the atoms
donate the lone pair electrons available Although theproperties of a coordinate covalent bond do not differ from
those of a normal covalent bond because all electrons are
a e no ma er w a e r source
2 Dative bond is usually applied for these few circumstances
below
a) To assist atom molecule ion that not yet achieved octet
configuration Making use of atom which has lone pair
electrons to those which are lack of electrons
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Sulphur dioxide SO2Sulphur trioxide SO3 Carbon monoxide CO
Ozone molecule O3 Water with hydrogen ion
S
OO
O
Ammonia with hydrogen ion Ammonia with boron trifluoride
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b) Formation of dimer - In order for some compounds which
have incomplete octet to achieve stability they tend to form
dimer or polymer among themselves by using dative bondTwo of the most common examples are aluminium trichloride
and beryllium dichloride
Monomer of AlCl3Dimer of aluminium chloride Al
2
Cl6
Monomer of BeCl2 Dimer of Be2Cl4 Polymer of (BeCl2)n
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c) Formation of coordination compounds - Coordination
compounds are substances that contain at least one complex
ion a species consisting of a central metal cation (either atransition metal or a main-group metal) that is bonded to
molecules andor anions called ligands via dative
(coordinative) bond
hexaaquacopper (II) ion [Cu(H2O)6]2+ tetraamminenickel (II) ion [Ni(NH3)4]
2+
Hexacyanoferrate (III) ion [Fe(CN)6]3- Trioxalatocobaltate(III)ion [Co(C2O4)3]
3-
3 2 3 Hybridisation Theory
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323 Hybridisation Theory
Valence bond theory
bull The basic principle of valence bond theory is that a covalentbond is formed when orbitals of two atoms overlap and the
overlapped region which is between the nuclei is occupied by
a pair of electrons The central themes of valence bond theory
derive from this principle
ndash Opposing spins of the electron pair ~ Stated in Paulis
orbitals has a maximum capacity of two electrons that musthave opposite spins For example when a covalent bond is
formed in molecule of hydrogen H2 the two 1s electrons of
two H atoms occupy the overlapping 1s orbitals and have
opposite spins
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bull Maximum overlap of bonding orbitals ~ The bond strength
depends on attraction of the nuclei for the shared electrons
so the greater the orbital overlap the stronger (morestable) the bond The extent of overlap depends on the
shapes and directions of the orbitals An s orbital is spherical
but p and d orbitals have more electron density in one
direction than in another Thus whenever possible a bond
involving p or d orbitals will be oriented in the direction that
maximizes overlap For example in hydrogen fluoride (HF)
bond the 1s orbital of H overlaps the half-filled 2p orbital of Falong the long axis of that orbital Any other direction would
result in less overlap and thus a weaker bond
H b idi ti f t i bit l T t f th
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bull Hybridisation of atomic orbitals ~ To account for the
bonding in simple diatomic molecules like HF we picture the
direct overlap of s and p orbitals of isolated atoms But howcan we account for the shapes of so many molecules and
polyatomic ions through the overlap of spherical s orbitals
dumbbell-shaped p orbitals and cloverleaf-shaped d orbitals
Linus Pauling proposed that the valence atomic orbitals in the
molecule are different from those in the isolated atoms The
spatial orientations of these new orbitals lead to more stable
bonds and are consistent with observed molecular shapesThe process of orbital mixing is called hybridisation and
the new atomic orbitals are called hybrid orbitals Two key
points about the number and type of hybrid orbitals are that
i The number of hybrid orbitals obtained equals the number ofatomic orbitals mixed
ii The type of hybrid orbitals obtained varies with the types of
atomic orbitals mixed
3 2 3 1 T f h b idi ti
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3231 Type of hybridisation
1 sp3 hybridisation ~ When four electron groups surround
the central atom the center atom involved must prepare fourorbitals with equal energies to overlap with the four
surrounding electron groups Valence Bond theory uses
hypothetical hybrid orbitals which are atomic orbitals
obtained when two or more non-equivalent orbitals of the
same atom combine in preparation for covalent bond
formation H bridisation is the term a lied to the mixin of
atomic orbitals in an atom (usually a central atom) togenerate a set of hybrid orbitals In the case of sp3 We can
generate four equivalent hybrid orbitals from the center
atom by mixing the s orbital and the three p orbitals
Explanation Energy level diagram Diagram of orbitals
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Explanation Energy level diagram Diagram of orbitals
a) Ground state
Carbon which act as thecenter atom has the valence
electron of 2s22p2
b) Excited state One of the electron from 2s
is promote to 2p orbitals -
equa energy eve
c) Hybridised state
One orbital of 2s and three
orbitals of 2p combined(hybrid) and rearrange
themselves to the shape
and orientation of a
tetrahedral shape
Molecular shape Tetrahedral
Angle between bond pair 10950
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2 sp2 hybridisation ~ When three electron groups surround
the central atom the center atom involved must prepare three
orbitals with equal energies to overlap with the three
surrounding electron groups In sp2 hybridisation three
equivalent (in terms of energy level) from the center atom by
mixing the one s orbital and the two p orbitals Using boron
trifluoride (BF3) as example sp2
hybridisation is explained
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8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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3 sp hybridisation ~ When two electron groups surround
the central atom we observe a linear shape which
means that the bonding orbitals must have a linear
orientation VB theory explains this by proposing thatmixing two nonequivalent orbitals of a central atom one
s and one p gives rise to two equivalent sp hybrid
orbitals that lie 1800 apart
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Beryllium which act as the
center atom has the valence
electron of 2s2
b) Excited state
One of the electron from 2s ispromote to 2p orbitals - equal
energy level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of a
linear shape
Molecular shape Linear
Angle between bond pair 1800
4 The concept of hybridisation is also useful to explain molecules with
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p y p
doubletriple bonds By using the concept of the direct overlapping
orbitals and side-touch lapping orbitals the formation of multiple
bonds in ethene C2H4 and ethyne C2H2 are described
a) Ethene C2H4 - Hybridisation take place for both carbon atoms in
ethene molecule is sp2 hybridisation
bull When both hybridised C is bonded together one of the hybridisedorbital overlapped directly between each other while the other two
hybridised orbitals overlapped directly with hydrogen atoms
z
atoms form a side-touch bond between each other and form anotherbond as shown in the diagram below
bull From the diagram with C=C there are two types of bond A sigma-
bond (σ-bond) is covalent bonds formed by orbitals overlapping end-
to-end with the electron density concentrated between the nuclei ofthe bonding atoms while the second type is called a pi bond (π-
bond) which is defined as a covalent bond formed by sideways
overlapping orbitals with electron density concentrated above and
below the plane of the nuclei of the bonding atoms
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital of 2s and two
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of atrigonal planar shape
Note that on unhybridised pz
orbital an electron is presence
Molecular shape Trigonal planar Angle between bond pair 1200
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b) Ethyne C H has the Lewis structure of H CequivC H which the
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b) Ethyne C2H2 has the Lewis structure of HminusCequivCminusH which the
bonding can be explain using sp hybridisation Table below
described how the hybridisation take place on each carbonatom and how the formation of triple bond occur
c) From the diagram formation of -CequivC- is due to the formation
of one sigma-bond by direct overlapping of one of the twohybrid orbitals on each C while the other two pi-bond bonds
are formed as a result of side-lapping of the each two
y z
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital from 2s and 2p
orbitals combined (hybrid) and
rearrange themselves to the
shape and orientation of a
linear shapeNote that on unhybridised py amp
pz orbitals electrons are
presence to form two
π-bonds Molecular shape Linear Angle between bond pair 1800
sp
pz py
5 Other examples of applications in valence bond theory includes the
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formation of nitrogen molecule N2 and hydrogen cyanide HCN
molecule
a) Nitrogen gas is Earths most abundant gas as it cover 78 of the
content of our air Nitrogen molecule is an inert gas thanks to its
short covalent bond and also its strong triple bond Therefore
a lot of heats are required to break the chemical bond of nitrogenbefore it can be applied in industries Using valence bond theory
the bonding of nitrogen molecule is explained in the diagram
b) So when the two hybridised nitrogen atom interacting amongeach other they hence formed a linear shape and two pi-bonds
(πminusbonds) are formed as a result of side-lapping of unhybridised
py and pz orbital respectively
N N
ExplanationEnergy level
diagramDiagram of orbitals
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diagram
a) Ground state
Both nitrogen whichact as the center atom
has the valence
electron of 2s22p3
b) Excited state One of the electron
from 2s is promote to 2p
orbitals - e ual ener
level
c) Hybridised state
One orbital of 2s and
one orbitals of 2p
combined (hybrid) and
rearrange themselves
to the shape and
orientation of a linear
shapeMolecular shape Linear
Angle between bond pair 1800
b) Whereas for hydrogen cyanide HCN both carbon atom and
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) y g y
nitrogen atom undergoes sp hybridisation in order to form a
linear structureExplanation
Energy level diagram for
carbon atom C
Energy level diagram for
nitrogen atom N
a) Ground state
Valence electronC 2s22p2
N 2s22p3
b) Excited state
One of the electron from
2s is promote to 2p
orbitals - equal energy
level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined
(hybrid) and form sp
hybridisation
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N
σσσσ
π
σσσσ
π
6 However there are a few limitation on valence bond theory such
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as when explaining the effect angle of bond-pair and bond-pair
electrons when there isare presence of lone pair electron in thecenter atom and also the difference of electronegativity
a) Bonding in ammonia NH3 and water H2O ~ Both NH3 and
H2O undergoes an arrangement similar to sp3 hybridisation
similar to that of C in methane Table below shows thehybridisation occur for nitrogen in ammonia and oxygen in water
Energy level diagram Energy level diagram
for nitrogen atom N for oxygen atom O
a) Ground state
Valence electron
N 2s22p3
O 2s2
2p4
b) Excited state
One of the electron
from 2s is promote to
2p orbitals
c) Hybridised state
One orbital of 2s and
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One orbital of 2s and
three orbitals from 2p
combined (hybrid) andrearrange themselves to
the shape and
orientation similar to that
of tetrahedral shape
Similar to arrangement in
tetrahedral
Similar to arrangement in
tetrahedral
of ammonia NH3 andwater H2O
Shape and angleShape trigonal pyramidal
Angle 1070
Shape bent
Angle 10450
Number of bond pair amp
lone pair electrons
Bond pair electron 3
Lone pair electron 1
Bond pair electron 2
Lone pair electron 2
bull The angle of bond pair - bond pair electrons in ammonia and0 0
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water are 1070 and 10450 respectively which is lesser than in
methane molecule (10950
) This can be explained by the factof the presence of lone-pair electrons in both ammonia and
water Since the lone pair - lone pair electron repulsion is
stronger than lone pair - bond pair electron repulsion than
bond pair - bond pair electron repulsion it is expected that therepulsion between lone pair - bond pair electrons in ammonia
is stronger than bond pair - bond pair electrons repulsion
hence caused the angle to squeeze to a smaller angle As
for water since there is a presence of lone pair - lone pair
electrons repulsion it results the bond pair - bond pair
electron repulsion to be much smaller hence caused the
angle to squeeze to a smaller angle Though valence bondtheory does not actually explain the hybridisation especially
when it relates to the repulsion occur involving lone pair
electron another theories shall be applied to study such
effect All these shall be discussed further in VSEPR theory
b) Effect of electronegativity and bonding angle between bond pair
- bond pair electrons ~
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- bond pair electrons
i As discussed earlier in the bonding of water (H2
O) the
molecular shape and angle is described in the diagram below
Sul hur which is also an element from Grou 16 formed h dro en
sulphide H2S when sulphur react with hydrogen However unlike
water the bonding angles is much smaller compare to water This is
due to the difference of electronegativity and also the bond length
between O and S in molecule Since O is more electronegative
compare to S the O-H bond is pulled closer toward O Furthermorethe bond length of O-H is shorter compare to S-H As a result the
bonding pair-bonding pair electrons repulse greater with each other in
H-O-H and caused the angle become greater
ii Another example is between NH3 and PH3 The orbital
f
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diagram for both NH3 and PH3 are described below Both
nitrogen N and phosphorous P are from the same groupwhich is Group 15
Similar to the above case P in phosphine is less
electronegative than N in ammonia and bond length of N-H
is shorter than P-H As a result H is pull closer to N and
repulsion between bonding of H-N-H is greater compare toH-P-H and caused the angle between H-N-H is greater
compare to H-P-H
324 Valence Shell Electron Pair Repulsion (VSEPR) Theory
1 M l l t i th th di i l t f t
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1 Molecular geometry is the three-dimensional arrangement of atoms
in a molecule A moleculersquos geometry affects its physical andchemical properties such as melting point boiling point density
and the types of reactions it undergoes The basic concept of
VSEPR theory is based on the three general rules below
a) As far as electron-pair repulsion is concerned double bonds andtriple bonds can be treated like single bonds This approximation is
good for qualitative purposes However you should realize that in
ldquo rdquo
because there are two or three bonds between two atoms theelectron density occupies more space
b) If a molecule has two or more resonance structures we can apply
the VSEPR model to any one of them Formal charges are usually
not shown
c) The order of repulsion strength of lone pair and bond pair are
lone-pair amp lone-pair electrons repulsion are the strongest
followed by lone-pair amp bond-pair electrons repulsion while
bond-pair amp bond-pair electrons repulsion is the weakest
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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Water H2O
H 1 s1 O 2 s22p4
Carbon dioxide CO2
C 2 s22p2 O 2 s22p4
Ammonia NH3
H 1 s1 N 2 s22p3Ethene C2H4
H 1 s1 C 2 s22p2
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Hydrogen cyanide HCN
H 1 s1 N 2 s22p3 C 2 s22p2
Ethanoic acid CH3COOH
H 1 s1 O 2 s22p4 C 2 s22p2
Tetrachloromethane CCl4C 2 s22p2 Cl 3s23p5
Ethyne C2H2
H 1 s1 C 2 s22p2
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a) Note that in ethene hydrogen cyanide and ethanoic acid all
the valence electrons are used in bonding there are no lone
pairs on the carbon atoms In fact most of the stablemolecules containing carbon do not have lone pairs on the
carbon atoms
b) Multiple bonds (double bond or triple bond) are shorter than single covalent bonds Bond length is defined as the
distance between the nuclei of two covalently bonded atoms
nitrogen triple bonds are shorter than double bonds whichin turn are shorter than single bonds The shorter multiple
bonds are also more stable than single bonds
c) Covalent bond not only exist in neutral molecule but also insome molecular ions Table below shows a few example of
molecular ions which have covalent bonds in its molecule
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Carbonate ion CO32-
C 2 s22p2 O 2 s22p4
Cyanide ion CN-
N 2 s22p3 C 2 s22p2
u p a e on 4-
S 3 s23p4 O 2 s22p4
ra e on 3-
N 2 s22p3 O 2 s22p4
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321 Exception of Octet Rules
1 From the Lewis structure sketched for sulphate ion SO42- we can
see that the center metal atom (sulphur) has more than 8electrons For this the molecule is described as molecule that can
expanded octet These molecules that have more than 8
electrons are located at Period 3 and below as these center
atoms have empty d-orbital available to expand the number ofelectrons positioned in the center atom Those center atom from
Period 2 such as C N O and F can only allocate a maximum
Example Both phosphorous and nitrogen are elements from Group15 with the valence electron of ns2np3 They can both form NCl3and PCl3 respectively when react with limited amount of chlorine
however under excess chlorine only PCl 5 can be formed but
not NCl 5 Explain the statement bolded
Solution This is due to phosphorous which is from period
3 have empty d-orbital to expand the octet However
nitrogen which is from Period 2 do not have emptyorbital and can only allocate 8 electrons in its shell
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Phosphorous pentachloride PCl5P 3 s23p3 Cl 3s23p5
Sulphur hexafluoride SF6
S 3 s23p4 F 2 s22p5
Bromine pentachloride BrCl5Cl 3s23p5 Br 4s24p5
Xenon tetrafluoride XeF4
Xe 5s25p6 F 2 s22p5
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2 There are also some stable covalent compounds which have
less than 8 electrons in their center atom (incomplete octet)
The center atoms are usually metals with great number ofvalence electrons with small atomic radius such as
beryllium boron and aluminium
BeCl2Be 2s2 Cl 3s23p5 BF3B 2 s22p1 F 2 s22p5 AlCl3Al 3s23p1 F 3 s23p5
These compound possessed stability due to their short bondlength between center atom with surround atom Furthermore
they can form resonance structure between the center atom
and surrounding atoms
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3 There are also some molecules which contain an odd
number of electrons Among the most common compounds
are nitrogen monoxide (NO) and nitrogen dioxide (NO2)
Odd-electron molecules are sometimes called radicals
Nitrogen monoxide NO
N 2 s22p3 O 2 s22p4
Nitrogen dioxide NO2
N 2 s22p3 O 2 s22p4
any ra ca s are g y reac ve e reason s t at t ere
is a tendency for the unpaired electron to form a
covalent bond with an unpaired electron on another
molecule For example when two nitrogen dioxide
molecules collide they form dinitrogen tetroxide
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322 Dative bond
1 As shown in the Lewis structure of nitrate ion one of the N-Obond is drawn as rarr The bond rarr placed is called as dative
bond (also known as coordinative bond) where dative
bond is defined as a covalent bond in which one of the atoms
donate the lone pair electrons available Although theproperties of a coordinate covalent bond do not differ from
those of a normal covalent bond because all electrons are
a e no ma er w a e r source
2 Dative bond is usually applied for these few circumstances
below
a) To assist atom molecule ion that not yet achieved octet
configuration Making use of atom which has lone pair
electrons to those which are lack of electrons
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Sulphur dioxide SO2Sulphur trioxide SO3 Carbon monoxide CO
Ozone molecule O3 Water with hydrogen ion
S
OO
O
Ammonia with hydrogen ion Ammonia with boron trifluoride
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b) Formation of dimer - In order for some compounds which
have incomplete octet to achieve stability they tend to form
dimer or polymer among themselves by using dative bondTwo of the most common examples are aluminium trichloride
and beryllium dichloride
Monomer of AlCl3Dimer of aluminium chloride Al
2
Cl6
Monomer of BeCl2 Dimer of Be2Cl4 Polymer of (BeCl2)n
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c) Formation of coordination compounds - Coordination
compounds are substances that contain at least one complex
ion a species consisting of a central metal cation (either atransition metal or a main-group metal) that is bonded to
molecules andor anions called ligands via dative
(coordinative) bond
hexaaquacopper (II) ion [Cu(H2O)6]2+ tetraamminenickel (II) ion [Ni(NH3)4]
2+
Hexacyanoferrate (III) ion [Fe(CN)6]3- Trioxalatocobaltate(III)ion [Co(C2O4)3]
3-
3 2 3 Hybridisation Theory
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323 Hybridisation Theory
Valence bond theory
bull The basic principle of valence bond theory is that a covalentbond is formed when orbitals of two atoms overlap and the
overlapped region which is between the nuclei is occupied by
a pair of electrons The central themes of valence bond theory
derive from this principle
ndash Opposing spins of the electron pair ~ Stated in Paulis
orbitals has a maximum capacity of two electrons that musthave opposite spins For example when a covalent bond is
formed in molecule of hydrogen H2 the two 1s electrons of
two H atoms occupy the overlapping 1s orbitals and have
opposite spins
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bull Maximum overlap of bonding orbitals ~ The bond strength
depends on attraction of the nuclei for the shared electrons
so the greater the orbital overlap the stronger (morestable) the bond The extent of overlap depends on the
shapes and directions of the orbitals An s orbital is spherical
but p and d orbitals have more electron density in one
direction than in another Thus whenever possible a bond
involving p or d orbitals will be oriented in the direction that
maximizes overlap For example in hydrogen fluoride (HF)
bond the 1s orbital of H overlaps the half-filled 2p orbital of Falong the long axis of that orbital Any other direction would
result in less overlap and thus a weaker bond
H b idi ti f t i bit l T t f th
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bull Hybridisation of atomic orbitals ~ To account for the
bonding in simple diatomic molecules like HF we picture the
direct overlap of s and p orbitals of isolated atoms But howcan we account for the shapes of so many molecules and
polyatomic ions through the overlap of spherical s orbitals
dumbbell-shaped p orbitals and cloverleaf-shaped d orbitals
Linus Pauling proposed that the valence atomic orbitals in the
molecule are different from those in the isolated atoms The
spatial orientations of these new orbitals lead to more stable
bonds and are consistent with observed molecular shapesThe process of orbital mixing is called hybridisation and
the new atomic orbitals are called hybrid orbitals Two key
points about the number and type of hybrid orbitals are that
i The number of hybrid orbitals obtained equals the number ofatomic orbitals mixed
ii The type of hybrid orbitals obtained varies with the types of
atomic orbitals mixed
3 2 3 1 T f h b idi ti
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3231 Type of hybridisation
1 sp3 hybridisation ~ When four electron groups surround
the central atom the center atom involved must prepare fourorbitals with equal energies to overlap with the four
surrounding electron groups Valence Bond theory uses
hypothetical hybrid orbitals which are atomic orbitals
obtained when two or more non-equivalent orbitals of the
same atom combine in preparation for covalent bond
formation H bridisation is the term a lied to the mixin of
atomic orbitals in an atom (usually a central atom) togenerate a set of hybrid orbitals In the case of sp3 We can
generate four equivalent hybrid orbitals from the center
atom by mixing the s orbital and the three p orbitals
Explanation Energy level diagram Diagram of orbitals
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Explanation Energy level diagram Diagram of orbitals
a) Ground state
Carbon which act as thecenter atom has the valence
electron of 2s22p2
b) Excited state One of the electron from 2s
is promote to 2p orbitals -
equa energy eve
c) Hybridised state
One orbital of 2s and three
orbitals of 2p combined(hybrid) and rearrange
themselves to the shape
and orientation of a
tetrahedral shape
Molecular shape Tetrahedral
Angle between bond pair 10950
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2 sp2 hybridisation ~ When three electron groups surround
the central atom the center atom involved must prepare three
orbitals with equal energies to overlap with the three
surrounding electron groups In sp2 hybridisation three
equivalent (in terms of energy level) from the center atom by
mixing the one s orbital and the two p orbitals Using boron
trifluoride (BF3) as example sp2
hybridisation is explained
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3 sp hybridisation ~ When two electron groups surround
the central atom we observe a linear shape which
means that the bonding orbitals must have a linear
orientation VB theory explains this by proposing thatmixing two nonequivalent orbitals of a central atom one
s and one p gives rise to two equivalent sp hybrid
orbitals that lie 1800 apart
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Beryllium which act as the
center atom has the valence
electron of 2s2
b) Excited state
One of the electron from 2s ispromote to 2p orbitals - equal
energy level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of a
linear shape
Molecular shape Linear
Angle between bond pair 1800
4 The concept of hybridisation is also useful to explain molecules with
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p y p
doubletriple bonds By using the concept of the direct overlapping
orbitals and side-touch lapping orbitals the formation of multiple
bonds in ethene C2H4 and ethyne C2H2 are described
a) Ethene C2H4 - Hybridisation take place for both carbon atoms in
ethene molecule is sp2 hybridisation
bull When both hybridised C is bonded together one of the hybridisedorbital overlapped directly between each other while the other two
hybridised orbitals overlapped directly with hydrogen atoms
z
atoms form a side-touch bond between each other and form anotherbond as shown in the diagram below
bull From the diagram with C=C there are two types of bond A sigma-
bond (σ-bond) is covalent bonds formed by orbitals overlapping end-
to-end with the electron density concentrated between the nuclei ofthe bonding atoms while the second type is called a pi bond (π-
bond) which is defined as a covalent bond formed by sideways
overlapping orbitals with electron density concentrated above and
below the plane of the nuclei of the bonding atoms
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital of 2s and two
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of atrigonal planar shape
Note that on unhybridised pz
orbital an electron is presence
Molecular shape Trigonal planar Angle between bond pair 1200
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b) Ethyne C H has the Lewis structure of H CequivC H which the
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b) Ethyne C2H2 has the Lewis structure of HminusCequivCminusH which the
bonding can be explain using sp hybridisation Table below
described how the hybridisation take place on each carbonatom and how the formation of triple bond occur
c) From the diagram formation of -CequivC- is due to the formation
of one sigma-bond by direct overlapping of one of the twohybrid orbitals on each C while the other two pi-bond bonds
are formed as a result of side-lapping of the each two
y z
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital from 2s and 2p
orbitals combined (hybrid) and
rearrange themselves to the
shape and orientation of a
linear shapeNote that on unhybridised py amp
pz orbitals electrons are
presence to form two
π-bonds Molecular shape Linear Angle between bond pair 1800
sp
pz py
5 Other examples of applications in valence bond theory includes the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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formation of nitrogen molecule N2 and hydrogen cyanide HCN
molecule
a) Nitrogen gas is Earths most abundant gas as it cover 78 of the
content of our air Nitrogen molecule is an inert gas thanks to its
short covalent bond and also its strong triple bond Therefore
a lot of heats are required to break the chemical bond of nitrogenbefore it can be applied in industries Using valence bond theory
the bonding of nitrogen molecule is explained in the diagram
b) So when the two hybridised nitrogen atom interacting amongeach other they hence formed a linear shape and two pi-bonds
(πminusbonds) are formed as a result of side-lapping of unhybridised
py and pz orbital respectively
N N
ExplanationEnergy level
diagramDiagram of orbitals
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diagram
a) Ground state
Both nitrogen whichact as the center atom
has the valence
electron of 2s22p3
b) Excited state One of the electron
from 2s is promote to 2p
orbitals - e ual ener
level
c) Hybridised state
One orbital of 2s and
one orbitals of 2p
combined (hybrid) and
rearrange themselves
to the shape and
orientation of a linear
shapeMolecular shape Linear
Angle between bond pair 1800
b) Whereas for hydrogen cyanide HCN both carbon atom and
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) y g y
nitrogen atom undergoes sp hybridisation in order to form a
linear structureExplanation
Energy level diagram for
carbon atom C
Energy level diagram for
nitrogen atom N
a) Ground state
Valence electronC 2s22p2
N 2s22p3
b) Excited state
One of the electron from
2s is promote to 2p
orbitals - equal energy
level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined
(hybrid) and form sp
hybridisation
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N
σσσσ
π
σσσσ
π
6 However there are a few limitation on valence bond theory such
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as when explaining the effect angle of bond-pair and bond-pair
electrons when there isare presence of lone pair electron in thecenter atom and also the difference of electronegativity
a) Bonding in ammonia NH3 and water H2O ~ Both NH3 and
H2O undergoes an arrangement similar to sp3 hybridisation
similar to that of C in methane Table below shows thehybridisation occur for nitrogen in ammonia and oxygen in water
Energy level diagram Energy level diagram
for nitrogen atom N for oxygen atom O
a) Ground state
Valence electron
N 2s22p3
O 2s2
2p4
b) Excited state
One of the electron
from 2s is promote to
2p orbitals
c) Hybridised state
One orbital of 2s and
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One orbital of 2s and
three orbitals from 2p
combined (hybrid) andrearrange themselves to
the shape and
orientation similar to that
of tetrahedral shape
Similar to arrangement in
tetrahedral
Similar to arrangement in
tetrahedral
of ammonia NH3 andwater H2O
Shape and angleShape trigonal pyramidal
Angle 1070
Shape bent
Angle 10450
Number of bond pair amp
lone pair electrons
Bond pair electron 3
Lone pair electron 1
Bond pair electron 2
Lone pair electron 2
bull The angle of bond pair - bond pair electrons in ammonia and0 0
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water are 1070 and 10450 respectively which is lesser than in
methane molecule (10950
) This can be explained by the factof the presence of lone-pair electrons in both ammonia and
water Since the lone pair - lone pair electron repulsion is
stronger than lone pair - bond pair electron repulsion than
bond pair - bond pair electron repulsion it is expected that therepulsion between lone pair - bond pair electrons in ammonia
is stronger than bond pair - bond pair electrons repulsion
hence caused the angle to squeeze to a smaller angle As
for water since there is a presence of lone pair - lone pair
electrons repulsion it results the bond pair - bond pair
electron repulsion to be much smaller hence caused the
angle to squeeze to a smaller angle Though valence bondtheory does not actually explain the hybridisation especially
when it relates to the repulsion occur involving lone pair
electron another theories shall be applied to study such
effect All these shall be discussed further in VSEPR theory
b) Effect of electronegativity and bonding angle between bond pair
- bond pair electrons ~
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- bond pair electrons
i As discussed earlier in the bonding of water (H2
O) the
molecular shape and angle is described in the diagram below
Sul hur which is also an element from Grou 16 formed h dro en
sulphide H2S when sulphur react with hydrogen However unlike
water the bonding angles is much smaller compare to water This is
due to the difference of electronegativity and also the bond length
between O and S in molecule Since O is more electronegative
compare to S the O-H bond is pulled closer toward O Furthermorethe bond length of O-H is shorter compare to S-H As a result the
bonding pair-bonding pair electrons repulse greater with each other in
H-O-H and caused the angle become greater
ii Another example is between NH3 and PH3 The orbital
f
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diagram for both NH3 and PH3 are described below Both
nitrogen N and phosphorous P are from the same groupwhich is Group 15
Similar to the above case P in phosphine is less
electronegative than N in ammonia and bond length of N-H
is shorter than P-H As a result H is pull closer to N and
repulsion between bonding of H-N-H is greater compare toH-P-H and caused the angle between H-N-H is greater
compare to H-P-H
324 Valence Shell Electron Pair Repulsion (VSEPR) Theory
1 M l l t i th th di i l t f t
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1 Molecular geometry is the three-dimensional arrangement of atoms
in a molecule A moleculersquos geometry affects its physical andchemical properties such as melting point boiling point density
and the types of reactions it undergoes The basic concept of
VSEPR theory is based on the three general rules below
a) As far as electron-pair repulsion is concerned double bonds andtriple bonds can be treated like single bonds This approximation is
good for qualitative purposes However you should realize that in
ldquo rdquo
because there are two or three bonds between two atoms theelectron density occupies more space
b) If a molecule has two or more resonance structures we can apply
the VSEPR model to any one of them Formal charges are usually
not shown
c) The order of repulsion strength of lone pair and bond pair are
lone-pair amp lone-pair electrons repulsion are the strongest
followed by lone-pair amp bond-pair electrons repulsion while
bond-pair amp bond-pair electrons repulsion is the weakest
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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Hydrogen cyanide HCN
H 1 s1 N 2 s22p3 C 2 s22p2
Ethanoic acid CH3COOH
H 1 s1 O 2 s22p4 C 2 s22p2
Tetrachloromethane CCl4C 2 s22p2 Cl 3s23p5
Ethyne C2H2
H 1 s1 C 2 s22p2
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a) Note that in ethene hydrogen cyanide and ethanoic acid all
the valence electrons are used in bonding there are no lone
pairs on the carbon atoms In fact most of the stablemolecules containing carbon do not have lone pairs on the
carbon atoms
b) Multiple bonds (double bond or triple bond) are shorter than single covalent bonds Bond length is defined as the
distance between the nuclei of two covalently bonded atoms
nitrogen triple bonds are shorter than double bonds whichin turn are shorter than single bonds The shorter multiple
bonds are also more stable than single bonds
c) Covalent bond not only exist in neutral molecule but also insome molecular ions Table below shows a few example of
molecular ions which have covalent bonds in its molecule
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Carbonate ion CO32-
C 2 s22p2 O 2 s22p4
Cyanide ion CN-
N 2 s22p3 C 2 s22p2
u p a e on 4-
S 3 s23p4 O 2 s22p4
ra e on 3-
N 2 s22p3 O 2 s22p4
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321 Exception of Octet Rules
1 From the Lewis structure sketched for sulphate ion SO42- we can
see that the center metal atom (sulphur) has more than 8electrons For this the molecule is described as molecule that can
expanded octet These molecules that have more than 8
electrons are located at Period 3 and below as these center
atoms have empty d-orbital available to expand the number ofelectrons positioned in the center atom Those center atom from
Period 2 such as C N O and F can only allocate a maximum
Example Both phosphorous and nitrogen are elements from Group15 with the valence electron of ns2np3 They can both form NCl3and PCl3 respectively when react with limited amount of chlorine
however under excess chlorine only PCl 5 can be formed but
not NCl 5 Explain the statement bolded
Solution This is due to phosphorous which is from period
3 have empty d-orbital to expand the octet However
nitrogen which is from Period 2 do not have emptyorbital and can only allocate 8 electrons in its shell
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Phosphorous pentachloride PCl5P 3 s23p3 Cl 3s23p5
Sulphur hexafluoride SF6
S 3 s23p4 F 2 s22p5
Bromine pentachloride BrCl5Cl 3s23p5 Br 4s24p5
Xenon tetrafluoride XeF4
Xe 5s25p6 F 2 s22p5
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2 There are also some stable covalent compounds which have
less than 8 electrons in their center atom (incomplete octet)
The center atoms are usually metals with great number ofvalence electrons with small atomic radius such as
beryllium boron and aluminium
BeCl2Be 2s2 Cl 3s23p5 BF3B 2 s22p1 F 2 s22p5 AlCl3Al 3s23p1 F 3 s23p5
These compound possessed stability due to their short bondlength between center atom with surround atom Furthermore
they can form resonance structure between the center atom
and surrounding atoms
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3 There are also some molecules which contain an odd
number of electrons Among the most common compounds
are nitrogen monoxide (NO) and nitrogen dioxide (NO2)
Odd-electron molecules are sometimes called radicals
Nitrogen monoxide NO
N 2 s22p3 O 2 s22p4
Nitrogen dioxide NO2
N 2 s22p3 O 2 s22p4
any ra ca s are g y reac ve e reason s t at t ere
is a tendency for the unpaired electron to form a
covalent bond with an unpaired electron on another
molecule For example when two nitrogen dioxide
molecules collide they form dinitrogen tetroxide
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322 Dative bond
1 As shown in the Lewis structure of nitrate ion one of the N-Obond is drawn as rarr The bond rarr placed is called as dative
bond (also known as coordinative bond) where dative
bond is defined as a covalent bond in which one of the atoms
donate the lone pair electrons available Although theproperties of a coordinate covalent bond do not differ from
those of a normal covalent bond because all electrons are
a e no ma er w a e r source
2 Dative bond is usually applied for these few circumstances
below
a) To assist atom molecule ion that not yet achieved octet
configuration Making use of atom which has lone pair
electrons to those which are lack of electrons
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Sulphur dioxide SO2Sulphur trioxide SO3 Carbon monoxide CO
Ozone molecule O3 Water with hydrogen ion
S
OO
O
Ammonia with hydrogen ion Ammonia with boron trifluoride
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b) Formation of dimer - In order for some compounds which
have incomplete octet to achieve stability they tend to form
dimer or polymer among themselves by using dative bondTwo of the most common examples are aluminium trichloride
and beryllium dichloride
Monomer of AlCl3Dimer of aluminium chloride Al
2
Cl6
Monomer of BeCl2 Dimer of Be2Cl4 Polymer of (BeCl2)n
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c) Formation of coordination compounds - Coordination
compounds are substances that contain at least one complex
ion a species consisting of a central metal cation (either atransition metal or a main-group metal) that is bonded to
molecules andor anions called ligands via dative
(coordinative) bond
hexaaquacopper (II) ion [Cu(H2O)6]2+ tetraamminenickel (II) ion [Ni(NH3)4]
2+
Hexacyanoferrate (III) ion [Fe(CN)6]3- Trioxalatocobaltate(III)ion [Co(C2O4)3]
3-
3 2 3 Hybridisation Theory
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323 Hybridisation Theory
Valence bond theory
bull The basic principle of valence bond theory is that a covalentbond is formed when orbitals of two atoms overlap and the
overlapped region which is between the nuclei is occupied by
a pair of electrons The central themes of valence bond theory
derive from this principle
ndash Opposing spins of the electron pair ~ Stated in Paulis
orbitals has a maximum capacity of two electrons that musthave opposite spins For example when a covalent bond is
formed in molecule of hydrogen H2 the two 1s electrons of
two H atoms occupy the overlapping 1s orbitals and have
opposite spins
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bull Maximum overlap of bonding orbitals ~ The bond strength
depends on attraction of the nuclei for the shared electrons
so the greater the orbital overlap the stronger (morestable) the bond The extent of overlap depends on the
shapes and directions of the orbitals An s orbital is spherical
but p and d orbitals have more electron density in one
direction than in another Thus whenever possible a bond
involving p or d orbitals will be oriented in the direction that
maximizes overlap For example in hydrogen fluoride (HF)
bond the 1s orbital of H overlaps the half-filled 2p orbital of Falong the long axis of that orbital Any other direction would
result in less overlap and thus a weaker bond
H b idi ti f t i bit l T t f th
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bull Hybridisation of atomic orbitals ~ To account for the
bonding in simple diatomic molecules like HF we picture the
direct overlap of s and p orbitals of isolated atoms But howcan we account for the shapes of so many molecules and
polyatomic ions through the overlap of spherical s orbitals
dumbbell-shaped p orbitals and cloverleaf-shaped d orbitals
Linus Pauling proposed that the valence atomic orbitals in the
molecule are different from those in the isolated atoms The
spatial orientations of these new orbitals lead to more stable
bonds and are consistent with observed molecular shapesThe process of orbital mixing is called hybridisation and
the new atomic orbitals are called hybrid orbitals Two key
points about the number and type of hybrid orbitals are that
i The number of hybrid orbitals obtained equals the number ofatomic orbitals mixed
ii The type of hybrid orbitals obtained varies with the types of
atomic orbitals mixed
3 2 3 1 T f h b idi ti
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3231 Type of hybridisation
1 sp3 hybridisation ~ When four electron groups surround
the central atom the center atom involved must prepare fourorbitals with equal energies to overlap with the four
surrounding electron groups Valence Bond theory uses
hypothetical hybrid orbitals which are atomic orbitals
obtained when two or more non-equivalent orbitals of the
same atom combine in preparation for covalent bond
formation H bridisation is the term a lied to the mixin of
atomic orbitals in an atom (usually a central atom) togenerate a set of hybrid orbitals In the case of sp3 We can
generate four equivalent hybrid orbitals from the center
atom by mixing the s orbital and the three p orbitals
Explanation Energy level diagram Diagram of orbitals
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Explanation Energy level diagram Diagram of orbitals
a) Ground state
Carbon which act as thecenter atom has the valence
electron of 2s22p2
b) Excited state One of the electron from 2s
is promote to 2p orbitals -
equa energy eve
c) Hybridised state
One orbital of 2s and three
orbitals of 2p combined(hybrid) and rearrange
themselves to the shape
and orientation of a
tetrahedral shape
Molecular shape Tetrahedral
Angle between bond pair 10950
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2 sp2 hybridisation ~ When three electron groups surround
the central atom the center atom involved must prepare three
orbitals with equal energies to overlap with the three
surrounding electron groups In sp2 hybridisation three
equivalent (in terms of energy level) from the center atom by
mixing the one s orbital and the two p orbitals Using boron
trifluoride (BF3) as example sp2
hybridisation is explained
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3 sp hybridisation ~ When two electron groups surround
the central atom we observe a linear shape which
means that the bonding orbitals must have a linear
orientation VB theory explains this by proposing thatmixing two nonequivalent orbitals of a central atom one
s and one p gives rise to two equivalent sp hybrid
orbitals that lie 1800 apart
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Beryllium which act as the
center atom has the valence
electron of 2s2
b) Excited state
One of the electron from 2s ispromote to 2p orbitals - equal
energy level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of a
linear shape
Molecular shape Linear
Angle between bond pair 1800
4 The concept of hybridisation is also useful to explain molecules with
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p y p
doubletriple bonds By using the concept of the direct overlapping
orbitals and side-touch lapping orbitals the formation of multiple
bonds in ethene C2H4 and ethyne C2H2 are described
a) Ethene C2H4 - Hybridisation take place for both carbon atoms in
ethene molecule is sp2 hybridisation
bull When both hybridised C is bonded together one of the hybridisedorbital overlapped directly between each other while the other two
hybridised orbitals overlapped directly with hydrogen atoms
z
atoms form a side-touch bond between each other and form anotherbond as shown in the diagram below
bull From the diagram with C=C there are two types of bond A sigma-
bond (σ-bond) is covalent bonds formed by orbitals overlapping end-
to-end with the electron density concentrated between the nuclei ofthe bonding atoms while the second type is called a pi bond (π-
bond) which is defined as a covalent bond formed by sideways
overlapping orbitals with electron density concentrated above and
below the plane of the nuclei of the bonding atoms
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital of 2s and two
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of atrigonal planar shape
Note that on unhybridised pz
orbital an electron is presence
Molecular shape Trigonal planar Angle between bond pair 1200
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b) Ethyne C H has the Lewis structure of H CequivC H which the
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b) Ethyne C2H2 has the Lewis structure of HminusCequivCminusH which the
bonding can be explain using sp hybridisation Table below
described how the hybridisation take place on each carbonatom and how the formation of triple bond occur
c) From the diagram formation of -CequivC- is due to the formation
of one sigma-bond by direct overlapping of one of the twohybrid orbitals on each C while the other two pi-bond bonds
are formed as a result of side-lapping of the each two
y z
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital from 2s and 2p
orbitals combined (hybrid) and
rearrange themselves to the
shape and orientation of a
linear shapeNote that on unhybridised py amp
pz orbitals electrons are
presence to form two
π-bonds Molecular shape Linear Angle between bond pair 1800
sp
pz py
5 Other examples of applications in valence bond theory includes the
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formation of nitrogen molecule N2 and hydrogen cyanide HCN
molecule
a) Nitrogen gas is Earths most abundant gas as it cover 78 of the
content of our air Nitrogen molecule is an inert gas thanks to its
short covalent bond and also its strong triple bond Therefore
a lot of heats are required to break the chemical bond of nitrogenbefore it can be applied in industries Using valence bond theory
the bonding of nitrogen molecule is explained in the diagram
b) So when the two hybridised nitrogen atom interacting amongeach other they hence formed a linear shape and two pi-bonds
(πminusbonds) are formed as a result of side-lapping of unhybridised
py and pz orbital respectively
N N
ExplanationEnergy level
diagramDiagram of orbitals
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diagram
a) Ground state
Both nitrogen whichact as the center atom
has the valence
electron of 2s22p3
b) Excited state One of the electron
from 2s is promote to 2p
orbitals - e ual ener
level
c) Hybridised state
One orbital of 2s and
one orbitals of 2p
combined (hybrid) and
rearrange themselves
to the shape and
orientation of a linear
shapeMolecular shape Linear
Angle between bond pair 1800
b) Whereas for hydrogen cyanide HCN both carbon atom and
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) y g y
nitrogen atom undergoes sp hybridisation in order to form a
linear structureExplanation
Energy level diagram for
carbon atom C
Energy level diagram for
nitrogen atom N
a) Ground state
Valence electronC 2s22p2
N 2s22p3
b) Excited state
One of the electron from
2s is promote to 2p
orbitals - equal energy
level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined
(hybrid) and form sp
hybridisation
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N
σσσσ
π
σσσσ
π
6 However there are a few limitation on valence bond theory such
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as when explaining the effect angle of bond-pair and bond-pair
electrons when there isare presence of lone pair electron in thecenter atom and also the difference of electronegativity
a) Bonding in ammonia NH3 and water H2O ~ Both NH3 and
H2O undergoes an arrangement similar to sp3 hybridisation
similar to that of C in methane Table below shows thehybridisation occur for nitrogen in ammonia and oxygen in water
Energy level diagram Energy level diagram
for nitrogen atom N for oxygen atom O
a) Ground state
Valence electron
N 2s22p3
O 2s2
2p4
b) Excited state
One of the electron
from 2s is promote to
2p orbitals
c) Hybridised state
One orbital of 2s and
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One orbital of 2s and
three orbitals from 2p
combined (hybrid) andrearrange themselves to
the shape and
orientation similar to that
of tetrahedral shape
Similar to arrangement in
tetrahedral
Similar to arrangement in
tetrahedral
of ammonia NH3 andwater H2O
Shape and angleShape trigonal pyramidal
Angle 1070
Shape bent
Angle 10450
Number of bond pair amp
lone pair electrons
Bond pair electron 3
Lone pair electron 1
Bond pair electron 2
Lone pair electron 2
bull The angle of bond pair - bond pair electrons in ammonia and0 0
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water are 1070 and 10450 respectively which is lesser than in
methane molecule (10950
) This can be explained by the factof the presence of lone-pair electrons in both ammonia and
water Since the lone pair - lone pair electron repulsion is
stronger than lone pair - bond pair electron repulsion than
bond pair - bond pair electron repulsion it is expected that therepulsion between lone pair - bond pair electrons in ammonia
is stronger than bond pair - bond pair electrons repulsion
hence caused the angle to squeeze to a smaller angle As
for water since there is a presence of lone pair - lone pair
electrons repulsion it results the bond pair - bond pair
electron repulsion to be much smaller hence caused the
angle to squeeze to a smaller angle Though valence bondtheory does not actually explain the hybridisation especially
when it relates to the repulsion occur involving lone pair
electron another theories shall be applied to study such
effect All these shall be discussed further in VSEPR theory
b) Effect of electronegativity and bonding angle between bond pair
- bond pair electrons ~
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- bond pair electrons
i As discussed earlier in the bonding of water (H2
O) the
molecular shape and angle is described in the diagram below
Sul hur which is also an element from Grou 16 formed h dro en
sulphide H2S when sulphur react with hydrogen However unlike
water the bonding angles is much smaller compare to water This is
due to the difference of electronegativity and also the bond length
between O and S in molecule Since O is more electronegative
compare to S the O-H bond is pulled closer toward O Furthermorethe bond length of O-H is shorter compare to S-H As a result the
bonding pair-bonding pair electrons repulse greater with each other in
H-O-H and caused the angle become greater
ii Another example is between NH3 and PH3 The orbital
f
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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diagram for both NH3 and PH3 are described below Both
nitrogen N and phosphorous P are from the same groupwhich is Group 15
Similar to the above case P in phosphine is less
electronegative than N in ammonia and bond length of N-H
is shorter than P-H As a result H is pull closer to N and
repulsion between bonding of H-N-H is greater compare toH-P-H and caused the angle between H-N-H is greater
compare to H-P-H
324 Valence Shell Electron Pair Repulsion (VSEPR) Theory
1 M l l t i th th di i l t f t
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1 Molecular geometry is the three-dimensional arrangement of atoms
in a molecule A moleculersquos geometry affects its physical andchemical properties such as melting point boiling point density
and the types of reactions it undergoes The basic concept of
VSEPR theory is based on the three general rules below
a) As far as electron-pair repulsion is concerned double bonds andtriple bonds can be treated like single bonds This approximation is
good for qualitative purposes However you should realize that in
ldquo rdquo
because there are two or three bonds between two atoms theelectron density occupies more space
b) If a molecule has two or more resonance structures we can apply
the VSEPR model to any one of them Formal charges are usually
not shown
c) The order of repulsion strength of lone pair and bond pair are
lone-pair amp lone-pair electrons repulsion are the strongest
followed by lone-pair amp bond-pair electrons repulsion while
bond-pair amp bond-pair electrons repulsion is the weakest
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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a) Note that in ethene hydrogen cyanide and ethanoic acid all
the valence electrons are used in bonding there are no lone
pairs on the carbon atoms In fact most of the stablemolecules containing carbon do not have lone pairs on the
carbon atoms
b) Multiple bonds (double bond or triple bond) are shorter than single covalent bonds Bond length is defined as the
distance between the nuclei of two covalently bonded atoms
nitrogen triple bonds are shorter than double bonds whichin turn are shorter than single bonds The shorter multiple
bonds are also more stable than single bonds
c) Covalent bond not only exist in neutral molecule but also insome molecular ions Table below shows a few example of
molecular ions which have covalent bonds in its molecule
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Carbonate ion CO32-
C 2 s22p2 O 2 s22p4
Cyanide ion CN-
N 2 s22p3 C 2 s22p2
u p a e on 4-
S 3 s23p4 O 2 s22p4
ra e on 3-
N 2 s22p3 O 2 s22p4
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321 Exception of Octet Rules
1 From the Lewis structure sketched for sulphate ion SO42- we can
see that the center metal atom (sulphur) has more than 8electrons For this the molecule is described as molecule that can
expanded octet These molecules that have more than 8
electrons are located at Period 3 and below as these center
atoms have empty d-orbital available to expand the number ofelectrons positioned in the center atom Those center atom from
Period 2 such as C N O and F can only allocate a maximum
Example Both phosphorous and nitrogen are elements from Group15 with the valence electron of ns2np3 They can both form NCl3and PCl3 respectively when react with limited amount of chlorine
however under excess chlorine only PCl 5 can be formed but
not NCl 5 Explain the statement bolded
Solution This is due to phosphorous which is from period
3 have empty d-orbital to expand the octet However
nitrogen which is from Period 2 do not have emptyorbital and can only allocate 8 electrons in its shell
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Phosphorous pentachloride PCl5P 3 s23p3 Cl 3s23p5
Sulphur hexafluoride SF6
S 3 s23p4 F 2 s22p5
Bromine pentachloride BrCl5Cl 3s23p5 Br 4s24p5
Xenon tetrafluoride XeF4
Xe 5s25p6 F 2 s22p5
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2 There are also some stable covalent compounds which have
less than 8 electrons in their center atom (incomplete octet)
The center atoms are usually metals with great number ofvalence electrons with small atomic radius such as
beryllium boron and aluminium
BeCl2Be 2s2 Cl 3s23p5 BF3B 2 s22p1 F 2 s22p5 AlCl3Al 3s23p1 F 3 s23p5
These compound possessed stability due to their short bondlength between center atom with surround atom Furthermore
they can form resonance structure between the center atom
and surrounding atoms
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3 There are also some molecules which contain an odd
number of electrons Among the most common compounds
are nitrogen monoxide (NO) and nitrogen dioxide (NO2)
Odd-electron molecules are sometimes called radicals
Nitrogen monoxide NO
N 2 s22p3 O 2 s22p4
Nitrogen dioxide NO2
N 2 s22p3 O 2 s22p4
any ra ca s are g y reac ve e reason s t at t ere
is a tendency for the unpaired electron to form a
covalent bond with an unpaired electron on another
molecule For example when two nitrogen dioxide
molecules collide they form dinitrogen tetroxide
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322 Dative bond
1 As shown in the Lewis structure of nitrate ion one of the N-Obond is drawn as rarr The bond rarr placed is called as dative
bond (also known as coordinative bond) where dative
bond is defined as a covalent bond in which one of the atoms
donate the lone pair electrons available Although theproperties of a coordinate covalent bond do not differ from
those of a normal covalent bond because all electrons are
a e no ma er w a e r source
2 Dative bond is usually applied for these few circumstances
below
a) To assist atom molecule ion that not yet achieved octet
configuration Making use of atom which has lone pair
electrons to those which are lack of electrons
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Sulphur dioxide SO2Sulphur trioxide SO3 Carbon monoxide CO
Ozone molecule O3 Water with hydrogen ion
S
OO
O
Ammonia with hydrogen ion Ammonia with boron trifluoride
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b) Formation of dimer - In order for some compounds which
have incomplete octet to achieve stability they tend to form
dimer or polymer among themselves by using dative bondTwo of the most common examples are aluminium trichloride
and beryllium dichloride
Monomer of AlCl3Dimer of aluminium chloride Al
2
Cl6
Monomer of BeCl2 Dimer of Be2Cl4 Polymer of (BeCl2)n
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c) Formation of coordination compounds - Coordination
compounds are substances that contain at least one complex
ion a species consisting of a central metal cation (either atransition metal or a main-group metal) that is bonded to
molecules andor anions called ligands via dative
(coordinative) bond
hexaaquacopper (II) ion [Cu(H2O)6]2+ tetraamminenickel (II) ion [Ni(NH3)4]
2+
Hexacyanoferrate (III) ion [Fe(CN)6]3- Trioxalatocobaltate(III)ion [Co(C2O4)3]
3-
3 2 3 Hybridisation Theory
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323 Hybridisation Theory
Valence bond theory
bull The basic principle of valence bond theory is that a covalentbond is formed when orbitals of two atoms overlap and the
overlapped region which is between the nuclei is occupied by
a pair of electrons The central themes of valence bond theory
derive from this principle
ndash Opposing spins of the electron pair ~ Stated in Paulis
orbitals has a maximum capacity of two electrons that musthave opposite spins For example when a covalent bond is
formed in molecule of hydrogen H2 the two 1s electrons of
two H atoms occupy the overlapping 1s orbitals and have
opposite spins
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bull Maximum overlap of bonding orbitals ~ The bond strength
depends on attraction of the nuclei for the shared electrons
so the greater the orbital overlap the stronger (morestable) the bond The extent of overlap depends on the
shapes and directions of the orbitals An s orbital is spherical
but p and d orbitals have more electron density in one
direction than in another Thus whenever possible a bond
involving p or d orbitals will be oriented in the direction that
maximizes overlap For example in hydrogen fluoride (HF)
bond the 1s orbital of H overlaps the half-filled 2p orbital of Falong the long axis of that orbital Any other direction would
result in less overlap and thus a weaker bond
H b idi ti f t i bit l T t f th
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bull Hybridisation of atomic orbitals ~ To account for the
bonding in simple diatomic molecules like HF we picture the
direct overlap of s and p orbitals of isolated atoms But howcan we account for the shapes of so many molecules and
polyatomic ions through the overlap of spherical s orbitals
dumbbell-shaped p orbitals and cloverleaf-shaped d orbitals
Linus Pauling proposed that the valence atomic orbitals in the
molecule are different from those in the isolated atoms The
spatial orientations of these new orbitals lead to more stable
bonds and are consistent with observed molecular shapesThe process of orbital mixing is called hybridisation and
the new atomic orbitals are called hybrid orbitals Two key
points about the number and type of hybrid orbitals are that
i The number of hybrid orbitals obtained equals the number ofatomic orbitals mixed
ii The type of hybrid orbitals obtained varies with the types of
atomic orbitals mixed
3 2 3 1 T f h b idi ti
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3231 Type of hybridisation
1 sp3 hybridisation ~ When four electron groups surround
the central atom the center atom involved must prepare fourorbitals with equal energies to overlap with the four
surrounding electron groups Valence Bond theory uses
hypothetical hybrid orbitals which are atomic orbitals
obtained when two or more non-equivalent orbitals of the
same atom combine in preparation for covalent bond
formation H bridisation is the term a lied to the mixin of
atomic orbitals in an atom (usually a central atom) togenerate a set of hybrid orbitals In the case of sp3 We can
generate four equivalent hybrid orbitals from the center
atom by mixing the s orbital and the three p orbitals
Explanation Energy level diagram Diagram of orbitals
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Explanation Energy level diagram Diagram of orbitals
a) Ground state
Carbon which act as thecenter atom has the valence
electron of 2s22p2
b) Excited state One of the electron from 2s
is promote to 2p orbitals -
equa energy eve
c) Hybridised state
One orbital of 2s and three
orbitals of 2p combined(hybrid) and rearrange
themselves to the shape
and orientation of a
tetrahedral shape
Molecular shape Tetrahedral
Angle between bond pair 10950
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2 sp2 hybridisation ~ When three electron groups surround
the central atom the center atom involved must prepare three
orbitals with equal energies to overlap with the three
surrounding electron groups In sp2 hybridisation three
equivalent (in terms of energy level) from the center atom by
mixing the one s orbital and the two p orbitals Using boron
trifluoride (BF3) as example sp2
hybridisation is explained
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3 sp hybridisation ~ When two electron groups surround
the central atom we observe a linear shape which
means that the bonding orbitals must have a linear
orientation VB theory explains this by proposing thatmixing two nonequivalent orbitals of a central atom one
s and one p gives rise to two equivalent sp hybrid
orbitals that lie 1800 apart
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Beryllium which act as the
center atom has the valence
electron of 2s2
b) Excited state
One of the electron from 2s ispromote to 2p orbitals - equal
energy level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of a
linear shape
Molecular shape Linear
Angle between bond pair 1800
4 The concept of hybridisation is also useful to explain molecules with
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p y p
doubletriple bonds By using the concept of the direct overlapping
orbitals and side-touch lapping orbitals the formation of multiple
bonds in ethene C2H4 and ethyne C2H2 are described
a) Ethene C2H4 - Hybridisation take place for both carbon atoms in
ethene molecule is sp2 hybridisation
bull When both hybridised C is bonded together one of the hybridisedorbital overlapped directly between each other while the other two
hybridised orbitals overlapped directly with hydrogen atoms
z
atoms form a side-touch bond between each other and form anotherbond as shown in the diagram below
bull From the diagram with C=C there are two types of bond A sigma-
bond (σ-bond) is covalent bonds formed by orbitals overlapping end-
to-end with the electron density concentrated between the nuclei ofthe bonding atoms while the second type is called a pi bond (π-
bond) which is defined as a covalent bond formed by sideways
overlapping orbitals with electron density concentrated above and
below the plane of the nuclei of the bonding atoms
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital of 2s and two
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of atrigonal planar shape
Note that on unhybridised pz
orbital an electron is presence
Molecular shape Trigonal planar Angle between bond pair 1200
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b) Ethyne C H has the Lewis structure of H CequivC H which the
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b) Ethyne C2H2 has the Lewis structure of HminusCequivCminusH which the
bonding can be explain using sp hybridisation Table below
described how the hybridisation take place on each carbonatom and how the formation of triple bond occur
c) From the diagram formation of -CequivC- is due to the formation
of one sigma-bond by direct overlapping of one of the twohybrid orbitals on each C while the other two pi-bond bonds
are formed as a result of side-lapping of the each two
y z
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital from 2s and 2p
orbitals combined (hybrid) and
rearrange themselves to the
shape and orientation of a
linear shapeNote that on unhybridised py amp
pz orbitals electrons are
presence to form two
π-bonds Molecular shape Linear Angle between bond pair 1800
sp
pz py
5 Other examples of applications in valence bond theory includes the
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formation of nitrogen molecule N2 and hydrogen cyanide HCN
molecule
a) Nitrogen gas is Earths most abundant gas as it cover 78 of the
content of our air Nitrogen molecule is an inert gas thanks to its
short covalent bond and also its strong triple bond Therefore
a lot of heats are required to break the chemical bond of nitrogenbefore it can be applied in industries Using valence bond theory
the bonding of nitrogen molecule is explained in the diagram
b) So when the two hybridised nitrogen atom interacting amongeach other they hence formed a linear shape and two pi-bonds
(πminusbonds) are formed as a result of side-lapping of unhybridised
py and pz orbital respectively
N N
ExplanationEnergy level
diagramDiagram of orbitals
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diagram
a) Ground state
Both nitrogen whichact as the center atom
has the valence
electron of 2s22p3
b) Excited state One of the electron
from 2s is promote to 2p
orbitals - e ual ener
level
c) Hybridised state
One orbital of 2s and
one orbitals of 2p
combined (hybrid) and
rearrange themselves
to the shape and
orientation of a linear
shapeMolecular shape Linear
Angle between bond pair 1800
b) Whereas for hydrogen cyanide HCN both carbon atom and
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) y g y
nitrogen atom undergoes sp hybridisation in order to form a
linear structureExplanation
Energy level diagram for
carbon atom C
Energy level diagram for
nitrogen atom N
a) Ground state
Valence electronC 2s22p2
N 2s22p3
b) Excited state
One of the electron from
2s is promote to 2p
orbitals - equal energy
level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined
(hybrid) and form sp
hybridisation
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N
σσσσ
π
σσσσ
π
6 However there are a few limitation on valence bond theory such
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as when explaining the effect angle of bond-pair and bond-pair
electrons when there isare presence of lone pair electron in thecenter atom and also the difference of electronegativity
a) Bonding in ammonia NH3 and water H2O ~ Both NH3 and
H2O undergoes an arrangement similar to sp3 hybridisation
similar to that of C in methane Table below shows thehybridisation occur for nitrogen in ammonia and oxygen in water
Energy level diagram Energy level diagram
for nitrogen atom N for oxygen atom O
a) Ground state
Valence electron
N 2s22p3
O 2s2
2p4
b) Excited state
One of the electron
from 2s is promote to
2p orbitals
c) Hybridised state
One orbital of 2s and
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One orbital of 2s and
three orbitals from 2p
combined (hybrid) andrearrange themselves to
the shape and
orientation similar to that
of tetrahedral shape
Similar to arrangement in
tetrahedral
Similar to arrangement in
tetrahedral
of ammonia NH3 andwater H2O
Shape and angleShape trigonal pyramidal
Angle 1070
Shape bent
Angle 10450
Number of bond pair amp
lone pair electrons
Bond pair electron 3
Lone pair electron 1
Bond pair electron 2
Lone pair electron 2
bull The angle of bond pair - bond pair electrons in ammonia and0 0
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water are 1070 and 10450 respectively which is lesser than in
methane molecule (10950
) This can be explained by the factof the presence of lone-pair electrons in both ammonia and
water Since the lone pair - lone pair electron repulsion is
stronger than lone pair - bond pair electron repulsion than
bond pair - bond pair electron repulsion it is expected that therepulsion between lone pair - bond pair electrons in ammonia
is stronger than bond pair - bond pair electrons repulsion
hence caused the angle to squeeze to a smaller angle As
for water since there is a presence of lone pair - lone pair
electrons repulsion it results the bond pair - bond pair
electron repulsion to be much smaller hence caused the
angle to squeeze to a smaller angle Though valence bondtheory does not actually explain the hybridisation especially
when it relates to the repulsion occur involving lone pair
electron another theories shall be applied to study such
effect All these shall be discussed further in VSEPR theory
b) Effect of electronegativity and bonding angle between bond pair
- bond pair electrons ~
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- bond pair electrons
i As discussed earlier in the bonding of water (H2
O) the
molecular shape and angle is described in the diagram below
Sul hur which is also an element from Grou 16 formed h dro en
sulphide H2S when sulphur react with hydrogen However unlike
water the bonding angles is much smaller compare to water This is
due to the difference of electronegativity and also the bond length
between O and S in molecule Since O is more electronegative
compare to S the O-H bond is pulled closer toward O Furthermorethe bond length of O-H is shorter compare to S-H As a result the
bonding pair-bonding pair electrons repulse greater with each other in
H-O-H and caused the angle become greater
ii Another example is between NH3 and PH3 The orbital
f
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diagram for both NH3 and PH3 are described below Both
nitrogen N and phosphorous P are from the same groupwhich is Group 15
Similar to the above case P in phosphine is less
electronegative than N in ammonia and bond length of N-H
is shorter than P-H As a result H is pull closer to N and
repulsion between bonding of H-N-H is greater compare toH-P-H and caused the angle between H-N-H is greater
compare to H-P-H
324 Valence Shell Electron Pair Repulsion (VSEPR) Theory
1 M l l t i th th di i l t f t
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1 Molecular geometry is the three-dimensional arrangement of atoms
in a molecule A moleculersquos geometry affects its physical andchemical properties such as melting point boiling point density
and the types of reactions it undergoes The basic concept of
VSEPR theory is based on the three general rules below
a) As far as electron-pair repulsion is concerned double bonds andtriple bonds can be treated like single bonds This approximation is
good for qualitative purposes However you should realize that in
ldquo rdquo
because there are two or three bonds between two atoms theelectron density occupies more space
b) If a molecule has two or more resonance structures we can apply
the VSEPR model to any one of them Formal charges are usually
not shown
c) The order of repulsion strength of lone pair and bond pair are
lone-pair amp lone-pair electrons repulsion are the strongest
followed by lone-pair amp bond-pair electrons repulsion while
bond-pair amp bond-pair electrons repulsion is the weakest
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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Carbonate ion CO32-
C 2 s22p2 O 2 s22p4
Cyanide ion CN-
N 2 s22p3 C 2 s22p2
u p a e on 4-
S 3 s23p4 O 2 s22p4
ra e on 3-
N 2 s22p3 O 2 s22p4
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321 Exception of Octet Rules
1 From the Lewis structure sketched for sulphate ion SO42- we can
see that the center metal atom (sulphur) has more than 8electrons For this the molecule is described as molecule that can
expanded octet These molecules that have more than 8
electrons are located at Period 3 and below as these center
atoms have empty d-orbital available to expand the number ofelectrons positioned in the center atom Those center atom from
Period 2 such as C N O and F can only allocate a maximum
Example Both phosphorous and nitrogen are elements from Group15 with the valence electron of ns2np3 They can both form NCl3and PCl3 respectively when react with limited amount of chlorine
however under excess chlorine only PCl 5 can be formed but
not NCl 5 Explain the statement bolded
Solution This is due to phosphorous which is from period
3 have empty d-orbital to expand the octet However
nitrogen which is from Period 2 do not have emptyorbital and can only allocate 8 electrons in its shell
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Phosphorous pentachloride PCl5P 3 s23p3 Cl 3s23p5
Sulphur hexafluoride SF6
S 3 s23p4 F 2 s22p5
Bromine pentachloride BrCl5Cl 3s23p5 Br 4s24p5
Xenon tetrafluoride XeF4
Xe 5s25p6 F 2 s22p5
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2 There are also some stable covalent compounds which have
less than 8 electrons in their center atom (incomplete octet)
The center atoms are usually metals with great number ofvalence electrons with small atomic radius such as
beryllium boron and aluminium
BeCl2Be 2s2 Cl 3s23p5 BF3B 2 s22p1 F 2 s22p5 AlCl3Al 3s23p1 F 3 s23p5
These compound possessed stability due to their short bondlength between center atom with surround atom Furthermore
they can form resonance structure between the center atom
and surrounding atoms
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3 There are also some molecules which contain an odd
number of electrons Among the most common compounds
are nitrogen monoxide (NO) and nitrogen dioxide (NO2)
Odd-electron molecules are sometimes called radicals
Nitrogen monoxide NO
N 2 s22p3 O 2 s22p4
Nitrogen dioxide NO2
N 2 s22p3 O 2 s22p4
any ra ca s are g y reac ve e reason s t at t ere
is a tendency for the unpaired electron to form a
covalent bond with an unpaired electron on another
molecule For example when two nitrogen dioxide
molecules collide they form dinitrogen tetroxide
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322 Dative bond
1 As shown in the Lewis structure of nitrate ion one of the N-Obond is drawn as rarr The bond rarr placed is called as dative
bond (also known as coordinative bond) where dative
bond is defined as a covalent bond in which one of the atoms
donate the lone pair electrons available Although theproperties of a coordinate covalent bond do not differ from
those of a normal covalent bond because all electrons are
a e no ma er w a e r source
2 Dative bond is usually applied for these few circumstances
below
a) To assist atom molecule ion that not yet achieved octet
configuration Making use of atom which has lone pair
electrons to those which are lack of electrons
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Sulphur dioxide SO2Sulphur trioxide SO3 Carbon monoxide CO
Ozone molecule O3 Water with hydrogen ion
S
OO
O
Ammonia with hydrogen ion Ammonia with boron trifluoride
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b) Formation of dimer - In order for some compounds which
have incomplete octet to achieve stability they tend to form
dimer or polymer among themselves by using dative bondTwo of the most common examples are aluminium trichloride
and beryllium dichloride
Monomer of AlCl3Dimer of aluminium chloride Al
2
Cl6
Monomer of BeCl2 Dimer of Be2Cl4 Polymer of (BeCl2)n
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c) Formation of coordination compounds - Coordination
compounds are substances that contain at least one complex
ion a species consisting of a central metal cation (either atransition metal or a main-group metal) that is bonded to
molecules andor anions called ligands via dative
(coordinative) bond
hexaaquacopper (II) ion [Cu(H2O)6]2+ tetraamminenickel (II) ion [Ni(NH3)4]
2+
Hexacyanoferrate (III) ion [Fe(CN)6]3- Trioxalatocobaltate(III)ion [Co(C2O4)3]
3-
3 2 3 Hybridisation Theory
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323 Hybridisation Theory
Valence bond theory
bull The basic principle of valence bond theory is that a covalentbond is formed when orbitals of two atoms overlap and the
overlapped region which is between the nuclei is occupied by
a pair of electrons The central themes of valence bond theory
derive from this principle
ndash Opposing spins of the electron pair ~ Stated in Paulis
orbitals has a maximum capacity of two electrons that musthave opposite spins For example when a covalent bond is
formed in molecule of hydrogen H2 the two 1s electrons of
two H atoms occupy the overlapping 1s orbitals and have
opposite spins
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bull Maximum overlap of bonding orbitals ~ The bond strength
depends on attraction of the nuclei for the shared electrons
so the greater the orbital overlap the stronger (morestable) the bond The extent of overlap depends on the
shapes and directions of the orbitals An s orbital is spherical
but p and d orbitals have more electron density in one
direction than in another Thus whenever possible a bond
involving p or d orbitals will be oriented in the direction that
maximizes overlap For example in hydrogen fluoride (HF)
bond the 1s orbital of H overlaps the half-filled 2p orbital of Falong the long axis of that orbital Any other direction would
result in less overlap and thus a weaker bond
H b idi ti f t i bit l T t f th
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bull Hybridisation of atomic orbitals ~ To account for the
bonding in simple diatomic molecules like HF we picture the
direct overlap of s and p orbitals of isolated atoms But howcan we account for the shapes of so many molecules and
polyatomic ions through the overlap of spherical s orbitals
dumbbell-shaped p orbitals and cloverleaf-shaped d orbitals
Linus Pauling proposed that the valence atomic orbitals in the
molecule are different from those in the isolated atoms The
spatial orientations of these new orbitals lead to more stable
bonds and are consistent with observed molecular shapesThe process of orbital mixing is called hybridisation and
the new atomic orbitals are called hybrid orbitals Two key
points about the number and type of hybrid orbitals are that
i The number of hybrid orbitals obtained equals the number ofatomic orbitals mixed
ii The type of hybrid orbitals obtained varies with the types of
atomic orbitals mixed
3 2 3 1 T f h b idi ti
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3231 Type of hybridisation
1 sp3 hybridisation ~ When four electron groups surround
the central atom the center atom involved must prepare fourorbitals with equal energies to overlap with the four
surrounding electron groups Valence Bond theory uses
hypothetical hybrid orbitals which are atomic orbitals
obtained when two or more non-equivalent orbitals of the
same atom combine in preparation for covalent bond
formation H bridisation is the term a lied to the mixin of
atomic orbitals in an atom (usually a central atom) togenerate a set of hybrid orbitals In the case of sp3 We can
generate four equivalent hybrid orbitals from the center
atom by mixing the s orbital and the three p orbitals
Explanation Energy level diagram Diagram of orbitals
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Explanation Energy level diagram Diagram of orbitals
a) Ground state
Carbon which act as thecenter atom has the valence
electron of 2s22p2
b) Excited state One of the electron from 2s
is promote to 2p orbitals -
equa energy eve
c) Hybridised state
One orbital of 2s and three
orbitals of 2p combined(hybrid) and rearrange
themselves to the shape
and orientation of a
tetrahedral shape
Molecular shape Tetrahedral
Angle between bond pair 10950
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2 sp2 hybridisation ~ When three electron groups surround
the central atom the center atom involved must prepare three
orbitals with equal energies to overlap with the three
surrounding electron groups In sp2 hybridisation three
equivalent (in terms of energy level) from the center atom by
mixing the one s orbital and the two p orbitals Using boron
trifluoride (BF3) as example sp2
hybridisation is explained
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3 sp hybridisation ~ When two electron groups surround
the central atom we observe a linear shape which
means that the bonding orbitals must have a linear
orientation VB theory explains this by proposing thatmixing two nonequivalent orbitals of a central atom one
s and one p gives rise to two equivalent sp hybrid
orbitals that lie 1800 apart
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Beryllium which act as the
center atom has the valence
electron of 2s2
b) Excited state
One of the electron from 2s ispromote to 2p orbitals - equal
energy level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of a
linear shape
Molecular shape Linear
Angle between bond pair 1800
4 The concept of hybridisation is also useful to explain molecules with
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p y p
doubletriple bonds By using the concept of the direct overlapping
orbitals and side-touch lapping orbitals the formation of multiple
bonds in ethene C2H4 and ethyne C2H2 are described
a) Ethene C2H4 - Hybridisation take place for both carbon atoms in
ethene molecule is sp2 hybridisation
bull When both hybridised C is bonded together one of the hybridisedorbital overlapped directly between each other while the other two
hybridised orbitals overlapped directly with hydrogen atoms
z
atoms form a side-touch bond between each other and form anotherbond as shown in the diagram below
bull From the diagram with C=C there are two types of bond A sigma-
bond (σ-bond) is covalent bonds formed by orbitals overlapping end-
to-end with the electron density concentrated between the nuclei ofthe bonding atoms while the second type is called a pi bond (π-
bond) which is defined as a covalent bond formed by sideways
overlapping orbitals with electron density concentrated above and
below the plane of the nuclei of the bonding atoms
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital of 2s and two
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of atrigonal planar shape
Note that on unhybridised pz
orbital an electron is presence
Molecular shape Trigonal planar Angle between bond pair 1200
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b) Ethyne C H has the Lewis structure of H CequivC H which the
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b) Ethyne C2H2 has the Lewis structure of HminusCequivCminusH which the
bonding can be explain using sp hybridisation Table below
described how the hybridisation take place on each carbonatom and how the formation of triple bond occur
c) From the diagram formation of -CequivC- is due to the formation
of one sigma-bond by direct overlapping of one of the twohybrid orbitals on each C while the other two pi-bond bonds
are formed as a result of side-lapping of the each two
y z
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital from 2s and 2p
orbitals combined (hybrid) and
rearrange themselves to the
shape and orientation of a
linear shapeNote that on unhybridised py amp
pz orbitals electrons are
presence to form two
π-bonds Molecular shape Linear Angle between bond pair 1800
sp
pz py
5 Other examples of applications in valence bond theory includes the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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formation of nitrogen molecule N2 and hydrogen cyanide HCN
molecule
a) Nitrogen gas is Earths most abundant gas as it cover 78 of the
content of our air Nitrogen molecule is an inert gas thanks to its
short covalent bond and also its strong triple bond Therefore
a lot of heats are required to break the chemical bond of nitrogenbefore it can be applied in industries Using valence bond theory
the bonding of nitrogen molecule is explained in the diagram
b) So when the two hybridised nitrogen atom interacting amongeach other they hence formed a linear shape and two pi-bonds
(πminusbonds) are formed as a result of side-lapping of unhybridised
py and pz orbital respectively
N N
ExplanationEnergy level
diagramDiagram of orbitals
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diagram
a) Ground state
Both nitrogen whichact as the center atom
has the valence
electron of 2s22p3
b) Excited state One of the electron
from 2s is promote to 2p
orbitals - e ual ener
level
c) Hybridised state
One orbital of 2s and
one orbitals of 2p
combined (hybrid) and
rearrange themselves
to the shape and
orientation of a linear
shapeMolecular shape Linear
Angle between bond pair 1800
b) Whereas for hydrogen cyanide HCN both carbon atom and
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) y g y
nitrogen atom undergoes sp hybridisation in order to form a
linear structureExplanation
Energy level diagram for
carbon atom C
Energy level diagram for
nitrogen atom N
a) Ground state
Valence electronC 2s22p2
N 2s22p3
b) Excited state
One of the electron from
2s is promote to 2p
orbitals - equal energy
level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined
(hybrid) and form sp
hybridisation
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N
σσσσ
π
σσσσ
π
6 However there are a few limitation on valence bond theory such
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as when explaining the effect angle of bond-pair and bond-pair
electrons when there isare presence of lone pair electron in thecenter atom and also the difference of electronegativity
a) Bonding in ammonia NH3 and water H2O ~ Both NH3 and
H2O undergoes an arrangement similar to sp3 hybridisation
similar to that of C in methane Table below shows thehybridisation occur for nitrogen in ammonia and oxygen in water
Energy level diagram Energy level diagram
for nitrogen atom N for oxygen atom O
a) Ground state
Valence electron
N 2s22p3
O 2s2
2p4
b) Excited state
One of the electron
from 2s is promote to
2p orbitals
c) Hybridised state
One orbital of 2s and
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One orbital of 2s and
three orbitals from 2p
combined (hybrid) andrearrange themselves to
the shape and
orientation similar to that
of tetrahedral shape
Similar to arrangement in
tetrahedral
Similar to arrangement in
tetrahedral
of ammonia NH3 andwater H2O
Shape and angleShape trigonal pyramidal
Angle 1070
Shape bent
Angle 10450
Number of bond pair amp
lone pair electrons
Bond pair electron 3
Lone pair electron 1
Bond pair electron 2
Lone pair electron 2
bull The angle of bond pair - bond pair electrons in ammonia and0 0
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water are 1070 and 10450 respectively which is lesser than in
methane molecule (10950
) This can be explained by the factof the presence of lone-pair electrons in both ammonia and
water Since the lone pair - lone pair electron repulsion is
stronger than lone pair - bond pair electron repulsion than
bond pair - bond pair electron repulsion it is expected that therepulsion between lone pair - bond pair electrons in ammonia
is stronger than bond pair - bond pair electrons repulsion
hence caused the angle to squeeze to a smaller angle As
for water since there is a presence of lone pair - lone pair
electrons repulsion it results the bond pair - bond pair
electron repulsion to be much smaller hence caused the
angle to squeeze to a smaller angle Though valence bondtheory does not actually explain the hybridisation especially
when it relates to the repulsion occur involving lone pair
electron another theories shall be applied to study such
effect All these shall be discussed further in VSEPR theory
b) Effect of electronegativity and bonding angle between bond pair
- bond pair electrons ~
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- bond pair electrons
i As discussed earlier in the bonding of water (H2
O) the
molecular shape and angle is described in the diagram below
Sul hur which is also an element from Grou 16 formed h dro en
sulphide H2S when sulphur react with hydrogen However unlike
water the bonding angles is much smaller compare to water This is
due to the difference of electronegativity and also the bond length
between O and S in molecule Since O is more electronegative
compare to S the O-H bond is pulled closer toward O Furthermorethe bond length of O-H is shorter compare to S-H As a result the
bonding pair-bonding pair electrons repulse greater with each other in
H-O-H and caused the angle become greater
ii Another example is between NH3 and PH3 The orbital
f
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diagram for both NH3 and PH3 are described below Both
nitrogen N and phosphorous P are from the same groupwhich is Group 15
Similar to the above case P in phosphine is less
electronegative than N in ammonia and bond length of N-H
is shorter than P-H As a result H is pull closer to N and
repulsion between bonding of H-N-H is greater compare toH-P-H and caused the angle between H-N-H is greater
compare to H-P-H
324 Valence Shell Electron Pair Repulsion (VSEPR) Theory
1 M l l t i th th di i l t f t
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1 Molecular geometry is the three-dimensional arrangement of atoms
in a molecule A moleculersquos geometry affects its physical andchemical properties such as melting point boiling point density
and the types of reactions it undergoes The basic concept of
VSEPR theory is based on the three general rules below
a) As far as electron-pair repulsion is concerned double bonds andtriple bonds can be treated like single bonds This approximation is
good for qualitative purposes However you should realize that in
ldquo rdquo
because there are two or three bonds between two atoms theelectron density occupies more space
b) If a molecule has two or more resonance structures we can apply
the VSEPR model to any one of them Formal charges are usually
not shown
c) The order of repulsion strength of lone pair and bond pair are
lone-pair amp lone-pair electrons repulsion are the strongest
followed by lone-pair amp bond-pair electrons repulsion while
bond-pair amp bond-pair electrons repulsion is the weakest
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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321 Exception of Octet Rules
1 From the Lewis structure sketched for sulphate ion SO42- we can
see that the center metal atom (sulphur) has more than 8electrons For this the molecule is described as molecule that can
expanded octet These molecules that have more than 8
electrons are located at Period 3 and below as these center
atoms have empty d-orbital available to expand the number ofelectrons positioned in the center atom Those center atom from
Period 2 such as C N O and F can only allocate a maximum
Example Both phosphorous and nitrogen are elements from Group15 with the valence electron of ns2np3 They can both form NCl3and PCl3 respectively when react with limited amount of chlorine
however under excess chlorine only PCl 5 can be formed but
not NCl 5 Explain the statement bolded
Solution This is due to phosphorous which is from period
3 have empty d-orbital to expand the octet However
nitrogen which is from Period 2 do not have emptyorbital and can only allocate 8 electrons in its shell
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Phosphorous pentachloride PCl5P 3 s23p3 Cl 3s23p5
Sulphur hexafluoride SF6
S 3 s23p4 F 2 s22p5
Bromine pentachloride BrCl5Cl 3s23p5 Br 4s24p5
Xenon tetrafluoride XeF4
Xe 5s25p6 F 2 s22p5
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2 There are also some stable covalent compounds which have
less than 8 electrons in their center atom (incomplete octet)
The center atoms are usually metals with great number ofvalence electrons with small atomic radius such as
beryllium boron and aluminium
BeCl2Be 2s2 Cl 3s23p5 BF3B 2 s22p1 F 2 s22p5 AlCl3Al 3s23p1 F 3 s23p5
These compound possessed stability due to their short bondlength between center atom with surround atom Furthermore
they can form resonance structure between the center atom
and surrounding atoms
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3 There are also some molecules which contain an odd
number of electrons Among the most common compounds
are nitrogen monoxide (NO) and nitrogen dioxide (NO2)
Odd-electron molecules are sometimes called radicals
Nitrogen monoxide NO
N 2 s22p3 O 2 s22p4
Nitrogen dioxide NO2
N 2 s22p3 O 2 s22p4
any ra ca s are g y reac ve e reason s t at t ere
is a tendency for the unpaired electron to form a
covalent bond with an unpaired electron on another
molecule For example when two nitrogen dioxide
molecules collide they form dinitrogen tetroxide
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322 Dative bond
1 As shown in the Lewis structure of nitrate ion one of the N-Obond is drawn as rarr The bond rarr placed is called as dative
bond (also known as coordinative bond) where dative
bond is defined as a covalent bond in which one of the atoms
donate the lone pair electrons available Although theproperties of a coordinate covalent bond do not differ from
those of a normal covalent bond because all electrons are
a e no ma er w a e r source
2 Dative bond is usually applied for these few circumstances
below
a) To assist atom molecule ion that not yet achieved octet
configuration Making use of atom which has lone pair
electrons to those which are lack of electrons
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Sulphur dioxide SO2Sulphur trioxide SO3 Carbon monoxide CO
Ozone molecule O3 Water with hydrogen ion
S
OO
O
Ammonia with hydrogen ion Ammonia with boron trifluoride
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b) Formation of dimer - In order for some compounds which
have incomplete octet to achieve stability they tend to form
dimer or polymer among themselves by using dative bondTwo of the most common examples are aluminium trichloride
and beryllium dichloride
Monomer of AlCl3Dimer of aluminium chloride Al
2
Cl6
Monomer of BeCl2 Dimer of Be2Cl4 Polymer of (BeCl2)n
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c) Formation of coordination compounds - Coordination
compounds are substances that contain at least one complex
ion a species consisting of a central metal cation (either atransition metal or a main-group metal) that is bonded to
molecules andor anions called ligands via dative
(coordinative) bond
hexaaquacopper (II) ion [Cu(H2O)6]2+ tetraamminenickel (II) ion [Ni(NH3)4]
2+
Hexacyanoferrate (III) ion [Fe(CN)6]3- Trioxalatocobaltate(III)ion [Co(C2O4)3]
3-
3 2 3 Hybridisation Theory
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323 Hybridisation Theory
Valence bond theory
bull The basic principle of valence bond theory is that a covalentbond is formed when orbitals of two atoms overlap and the
overlapped region which is between the nuclei is occupied by
a pair of electrons The central themes of valence bond theory
derive from this principle
ndash Opposing spins of the electron pair ~ Stated in Paulis
orbitals has a maximum capacity of two electrons that musthave opposite spins For example when a covalent bond is
formed in molecule of hydrogen H2 the two 1s electrons of
two H atoms occupy the overlapping 1s orbitals and have
opposite spins
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bull Maximum overlap of bonding orbitals ~ The bond strength
depends on attraction of the nuclei for the shared electrons
so the greater the orbital overlap the stronger (morestable) the bond The extent of overlap depends on the
shapes and directions of the orbitals An s orbital is spherical
but p and d orbitals have more electron density in one
direction than in another Thus whenever possible a bond
involving p or d orbitals will be oriented in the direction that
maximizes overlap For example in hydrogen fluoride (HF)
bond the 1s orbital of H overlaps the half-filled 2p orbital of Falong the long axis of that orbital Any other direction would
result in less overlap and thus a weaker bond
H b idi ti f t i bit l T t f th
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bull Hybridisation of atomic orbitals ~ To account for the
bonding in simple diatomic molecules like HF we picture the
direct overlap of s and p orbitals of isolated atoms But howcan we account for the shapes of so many molecules and
polyatomic ions through the overlap of spherical s orbitals
dumbbell-shaped p orbitals and cloverleaf-shaped d orbitals
Linus Pauling proposed that the valence atomic orbitals in the
molecule are different from those in the isolated atoms The
spatial orientations of these new orbitals lead to more stable
bonds and are consistent with observed molecular shapesThe process of orbital mixing is called hybridisation and
the new atomic orbitals are called hybrid orbitals Two key
points about the number and type of hybrid orbitals are that
i The number of hybrid orbitals obtained equals the number ofatomic orbitals mixed
ii The type of hybrid orbitals obtained varies with the types of
atomic orbitals mixed
3 2 3 1 T f h b idi ti
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3231 Type of hybridisation
1 sp3 hybridisation ~ When four electron groups surround
the central atom the center atom involved must prepare fourorbitals with equal energies to overlap with the four
surrounding electron groups Valence Bond theory uses
hypothetical hybrid orbitals which are atomic orbitals
obtained when two or more non-equivalent orbitals of the
same atom combine in preparation for covalent bond
formation H bridisation is the term a lied to the mixin of
atomic orbitals in an atom (usually a central atom) togenerate a set of hybrid orbitals In the case of sp3 We can
generate four equivalent hybrid orbitals from the center
atom by mixing the s orbital and the three p orbitals
Explanation Energy level diagram Diagram of orbitals
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Explanation Energy level diagram Diagram of orbitals
a) Ground state
Carbon which act as thecenter atom has the valence
electron of 2s22p2
b) Excited state One of the electron from 2s
is promote to 2p orbitals -
equa energy eve
c) Hybridised state
One orbital of 2s and three
orbitals of 2p combined(hybrid) and rearrange
themselves to the shape
and orientation of a
tetrahedral shape
Molecular shape Tetrahedral
Angle between bond pair 10950
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2 sp2 hybridisation ~ When three electron groups surround
the central atom the center atom involved must prepare three
orbitals with equal energies to overlap with the three
surrounding electron groups In sp2 hybridisation three
equivalent (in terms of energy level) from the center atom by
mixing the one s orbital and the two p orbitals Using boron
trifluoride (BF3) as example sp2
hybridisation is explained
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3 sp hybridisation ~ When two electron groups surround
the central atom we observe a linear shape which
means that the bonding orbitals must have a linear
orientation VB theory explains this by proposing thatmixing two nonequivalent orbitals of a central atom one
s and one p gives rise to two equivalent sp hybrid
orbitals that lie 1800 apart
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Beryllium which act as the
center atom has the valence
electron of 2s2
b) Excited state
One of the electron from 2s ispromote to 2p orbitals - equal
energy level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of a
linear shape
Molecular shape Linear
Angle between bond pair 1800
4 The concept of hybridisation is also useful to explain molecules with
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p y p
doubletriple bonds By using the concept of the direct overlapping
orbitals and side-touch lapping orbitals the formation of multiple
bonds in ethene C2H4 and ethyne C2H2 are described
a) Ethene C2H4 - Hybridisation take place for both carbon atoms in
ethene molecule is sp2 hybridisation
bull When both hybridised C is bonded together one of the hybridisedorbital overlapped directly between each other while the other two
hybridised orbitals overlapped directly with hydrogen atoms
z
atoms form a side-touch bond between each other and form anotherbond as shown in the diagram below
bull From the diagram with C=C there are two types of bond A sigma-
bond (σ-bond) is covalent bonds formed by orbitals overlapping end-
to-end with the electron density concentrated between the nuclei ofthe bonding atoms while the second type is called a pi bond (π-
bond) which is defined as a covalent bond formed by sideways
overlapping orbitals with electron density concentrated above and
below the plane of the nuclei of the bonding atoms
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital of 2s and two
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of atrigonal planar shape
Note that on unhybridised pz
orbital an electron is presence
Molecular shape Trigonal planar Angle between bond pair 1200
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b) Ethyne C H has the Lewis structure of H CequivC H which the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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b) Ethyne C2H2 has the Lewis structure of HminusCequivCminusH which the
bonding can be explain using sp hybridisation Table below
described how the hybridisation take place on each carbonatom and how the formation of triple bond occur
c) From the diagram formation of -CequivC- is due to the formation
of one sigma-bond by direct overlapping of one of the twohybrid orbitals on each C while the other two pi-bond bonds
are formed as a result of side-lapping of the each two
y z
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital from 2s and 2p
orbitals combined (hybrid) and
rearrange themselves to the
shape and orientation of a
linear shapeNote that on unhybridised py amp
pz orbitals electrons are
presence to form two
π-bonds Molecular shape Linear Angle between bond pair 1800
sp
pz py
5 Other examples of applications in valence bond theory includes the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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formation of nitrogen molecule N2 and hydrogen cyanide HCN
molecule
a) Nitrogen gas is Earths most abundant gas as it cover 78 of the
content of our air Nitrogen molecule is an inert gas thanks to its
short covalent bond and also its strong triple bond Therefore
a lot of heats are required to break the chemical bond of nitrogenbefore it can be applied in industries Using valence bond theory
the bonding of nitrogen molecule is explained in the diagram
b) So when the two hybridised nitrogen atom interacting amongeach other they hence formed a linear shape and two pi-bonds
(πminusbonds) are formed as a result of side-lapping of unhybridised
py and pz orbital respectively
N N
ExplanationEnergy level
diagramDiagram of orbitals
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diagram
a) Ground state
Both nitrogen whichact as the center atom
has the valence
electron of 2s22p3
b) Excited state One of the electron
from 2s is promote to 2p
orbitals - e ual ener
level
c) Hybridised state
One orbital of 2s and
one orbitals of 2p
combined (hybrid) and
rearrange themselves
to the shape and
orientation of a linear
shapeMolecular shape Linear
Angle between bond pair 1800
b) Whereas for hydrogen cyanide HCN both carbon atom and
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) y g y
nitrogen atom undergoes sp hybridisation in order to form a
linear structureExplanation
Energy level diagram for
carbon atom C
Energy level diagram for
nitrogen atom N
a) Ground state
Valence electronC 2s22p2
N 2s22p3
b) Excited state
One of the electron from
2s is promote to 2p
orbitals - equal energy
level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined
(hybrid) and form sp
hybridisation
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N
σσσσ
π
σσσσ
π
6 However there are a few limitation on valence bond theory such
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as when explaining the effect angle of bond-pair and bond-pair
electrons when there isare presence of lone pair electron in thecenter atom and also the difference of electronegativity
a) Bonding in ammonia NH3 and water H2O ~ Both NH3 and
H2O undergoes an arrangement similar to sp3 hybridisation
similar to that of C in methane Table below shows thehybridisation occur for nitrogen in ammonia and oxygen in water
Energy level diagram Energy level diagram
for nitrogen atom N for oxygen atom O
a) Ground state
Valence electron
N 2s22p3
O 2s2
2p4
b) Excited state
One of the electron
from 2s is promote to
2p orbitals
c) Hybridised state
One orbital of 2s and
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One orbital of 2s and
three orbitals from 2p
combined (hybrid) andrearrange themselves to
the shape and
orientation similar to that
of tetrahedral shape
Similar to arrangement in
tetrahedral
Similar to arrangement in
tetrahedral
of ammonia NH3 andwater H2O
Shape and angleShape trigonal pyramidal
Angle 1070
Shape bent
Angle 10450
Number of bond pair amp
lone pair electrons
Bond pair electron 3
Lone pair electron 1
Bond pair electron 2
Lone pair electron 2
bull The angle of bond pair - bond pair electrons in ammonia and0 0
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water are 1070 and 10450 respectively which is lesser than in
methane molecule (10950
) This can be explained by the factof the presence of lone-pair electrons in both ammonia and
water Since the lone pair - lone pair electron repulsion is
stronger than lone pair - bond pair electron repulsion than
bond pair - bond pair electron repulsion it is expected that therepulsion between lone pair - bond pair electrons in ammonia
is stronger than bond pair - bond pair electrons repulsion
hence caused the angle to squeeze to a smaller angle As
for water since there is a presence of lone pair - lone pair
electrons repulsion it results the bond pair - bond pair
electron repulsion to be much smaller hence caused the
angle to squeeze to a smaller angle Though valence bondtheory does not actually explain the hybridisation especially
when it relates to the repulsion occur involving lone pair
electron another theories shall be applied to study such
effect All these shall be discussed further in VSEPR theory
b) Effect of electronegativity and bonding angle between bond pair
- bond pair electrons ~
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- bond pair electrons
i As discussed earlier in the bonding of water (H2
O) the
molecular shape and angle is described in the diagram below
Sul hur which is also an element from Grou 16 formed h dro en
sulphide H2S when sulphur react with hydrogen However unlike
water the bonding angles is much smaller compare to water This is
due to the difference of electronegativity and also the bond length
between O and S in molecule Since O is more electronegative
compare to S the O-H bond is pulled closer toward O Furthermorethe bond length of O-H is shorter compare to S-H As a result the
bonding pair-bonding pair electrons repulse greater with each other in
H-O-H and caused the angle become greater
ii Another example is between NH3 and PH3 The orbital
f
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diagram for both NH3 and PH3 are described below Both
nitrogen N and phosphorous P are from the same groupwhich is Group 15
Similar to the above case P in phosphine is less
electronegative than N in ammonia and bond length of N-H
is shorter than P-H As a result H is pull closer to N and
repulsion between bonding of H-N-H is greater compare toH-P-H and caused the angle between H-N-H is greater
compare to H-P-H
324 Valence Shell Electron Pair Repulsion (VSEPR) Theory
1 M l l t i th th di i l t f t
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1 Molecular geometry is the three-dimensional arrangement of atoms
in a molecule A moleculersquos geometry affects its physical andchemical properties such as melting point boiling point density
and the types of reactions it undergoes The basic concept of
VSEPR theory is based on the three general rules below
a) As far as electron-pair repulsion is concerned double bonds andtriple bonds can be treated like single bonds This approximation is
good for qualitative purposes However you should realize that in
ldquo rdquo
because there are two or three bonds between two atoms theelectron density occupies more space
b) If a molecule has two or more resonance structures we can apply
the VSEPR model to any one of them Formal charges are usually
not shown
c) The order of repulsion strength of lone pair and bond pair are
lone-pair amp lone-pair electrons repulsion are the strongest
followed by lone-pair amp bond-pair electrons repulsion while
bond-pair amp bond-pair electrons repulsion is the weakest
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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Phosphorous pentachloride PCl5P 3 s23p3 Cl 3s23p5
Sulphur hexafluoride SF6
S 3 s23p4 F 2 s22p5
Bromine pentachloride BrCl5Cl 3s23p5 Br 4s24p5
Xenon tetrafluoride XeF4
Xe 5s25p6 F 2 s22p5
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2 There are also some stable covalent compounds which have
less than 8 electrons in their center atom (incomplete octet)
The center atoms are usually metals with great number ofvalence electrons with small atomic radius such as
beryllium boron and aluminium
BeCl2Be 2s2 Cl 3s23p5 BF3B 2 s22p1 F 2 s22p5 AlCl3Al 3s23p1 F 3 s23p5
These compound possessed stability due to their short bondlength between center atom with surround atom Furthermore
they can form resonance structure between the center atom
and surrounding atoms
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3 There are also some molecules which contain an odd
number of electrons Among the most common compounds
are nitrogen monoxide (NO) and nitrogen dioxide (NO2)
Odd-electron molecules are sometimes called radicals
Nitrogen monoxide NO
N 2 s22p3 O 2 s22p4
Nitrogen dioxide NO2
N 2 s22p3 O 2 s22p4
any ra ca s are g y reac ve e reason s t at t ere
is a tendency for the unpaired electron to form a
covalent bond with an unpaired electron on another
molecule For example when two nitrogen dioxide
molecules collide they form dinitrogen tetroxide
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322 Dative bond
1 As shown in the Lewis structure of nitrate ion one of the N-Obond is drawn as rarr The bond rarr placed is called as dative
bond (also known as coordinative bond) where dative
bond is defined as a covalent bond in which one of the atoms
donate the lone pair electrons available Although theproperties of a coordinate covalent bond do not differ from
those of a normal covalent bond because all electrons are
a e no ma er w a e r source
2 Dative bond is usually applied for these few circumstances
below
a) To assist atom molecule ion that not yet achieved octet
configuration Making use of atom which has lone pair
electrons to those which are lack of electrons
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Sulphur dioxide SO2Sulphur trioxide SO3 Carbon monoxide CO
Ozone molecule O3 Water with hydrogen ion
S
OO
O
Ammonia with hydrogen ion Ammonia with boron trifluoride
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b) Formation of dimer - In order for some compounds which
have incomplete octet to achieve stability they tend to form
dimer or polymer among themselves by using dative bondTwo of the most common examples are aluminium trichloride
and beryllium dichloride
Monomer of AlCl3Dimer of aluminium chloride Al
2
Cl6
Monomer of BeCl2 Dimer of Be2Cl4 Polymer of (BeCl2)n
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c) Formation of coordination compounds - Coordination
compounds are substances that contain at least one complex
ion a species consisting of a central metal cation (either atransition metal or a main-group metal) that is bonded to
molecules andor anions called ligands via dative
(coordinative) bond
hexaaquacopper (II) ion [Cu(H2O)6]2+ tetraamminenickel (II) ion [Ni(NH3)4]
2+
Hexacyanoferrate (III) ion [Fe(CN)6]3- Trioxalatocobaltate(III)ion [Co(C2O4)3]
3-
3 2 3 Hybridisation Theory
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323 Hybridisation Theory
Valence bond theory
bull The basic principle of valence bond theory is that a covalentbond is formed when orbitals of two atoms overlap and the
overlapped region which is between the nuclei is occupied by
a pair of electrons The central themes of valence bond theory
derive from this principle
ndash Opposing spins of the electron pair ~ Stated in Paulis
orbitals has a maximum capacity of two electrons that musthave opposite spins For example when a covalent bond is
formed in molecule of hydrogen H2 the two 1s electrons of
two H atoms occupy the overlapping 1s orbitals and have
opposite spins
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bull Maximum overlap of bonding orbitals ~ The bond strength
depends on attraction of the nuclei for the shared electrons
so the greater the orbital overlap the stronger (morestable) the bond The extent of overlap depends on the
shapes and directions of the orbitals An s orbital is spherical
but p and d orbitals have more electron density in one
direction than in another Thus whenever possible a bond
involving p or d orbitals will be oriented in the direction that
maximizes overlap For example in hydrogen fluoride (HF)
bond the 1s orbital of H overlaps the half-filled 2p orbital of Falong the long axis of that orbital Any other direction would
result in less overlap and thus a weaker bond
H b idi ti f t i bit l T t f th
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bull Hybridisation of atomic orbitals ~ To account for the
bonding in simple diatomic molecules like HF we picture the
direct overlap of s and p orbitals of isolated atoms But howcan we account for the shapes of so many molecules and
polyatomic ions through the overlap of spherical s orbitals
dumbbell-shaped p orbitals and cloverleaf-shaped d orbitals
Linus Pauling proposed that the valence atomic orbitals in the
molecule are different from those in the isolated atoms The
spatial orientations of these new orbitals lead to more stable
bonds and are consistent with observed molecular shapesThe process of orbital mixing is called hybridisation and
the new atomic orbitals are called hybrid orbitals Two key
points about the number and type of hybrid orbitals are that
i The number of hybrid orbitals obtained equals the number ofatomic orbitals mixed
ii The type of hybrid orbitals obtained varies with the types of
atomic orbitals mixed
3 2 3 1 T f h b idi ti
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3231 Type of hybridisation
1 sp3 hybridisation ~ When four electron groups surround
the central atom the center atom involved must prepare fourorbitals with equal energies to overlap with the four
surrounding electron groups Valence Bond theory uses
hypothetical hybrid orbitals which are atomic orbitals
obtained when two or more non-equivalent orbitals of the
same atom combine in preparation for covalent bond
formation H bridisation is the term a lied to the mixin of
atomic orbitals in an atom (usually a central atom) togenerate a set of hybrid orbitals In the case of sp3 We can
generate four equivalent hybrid orbitals from the center
atom by mixing the s orbital and the three p orbitals
Explanation Energy level diagram Diagram of orbitals
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Explanation Energy level diagram Diagram of orbitals
a) Ground state
Carbon which act as thecenter atom has the valence
electron of 2s22p2
b) Excited state One of the electron from 2s
is promote to 2p orbitals -
equa energy eve
c) Hybridised state
One orbital of 2s and three
orbitals of 2p combined(hybrid) and rearrange
themselves to the shape
and orientation of a
tetrahedral shape
Molecular shape Tetrahedral
Angle between bond pair 10950
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2 sp2 hybridisation ~ When three electron groups surround
the central atom the center atom involved must prepare three
orbitals with equal energies to overlap with the three
surrounding electron groups In sp2 hybridisation three
equivalent (in terms of energy level) from the center atom by
mixing the one s orbital and the two p orbitals Using boron
trifluoride (BF3) as example sp2
hybridisation is explained
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3 sp hybridisation ~ When two electron groups surround
the central atom we observe a linear shape which
means that the bonding orbitals must have a linear
orientation VB theory explains this by proposing thatmixing two nonequivalent orbitals of a central atom one
s and one p gives rise to two equivalent sp hybrid
orbitals that lie 1800 apart
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Beryllium which act as the
center atom has the valence
electron of 2s2
b) Excited state
One of the electron from 2s ispromote to 2p orbitals - equal
energy level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of a
linear shape
Molecular shape Linear
Angle between bond pair 1800
4 The concept of hybridisation is also useful to explain molecules with
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p y p
doubletriple bonds By using the concept of the direct overlapping
orbitals and side-touch lapping orbitals the formation of multiple
bonds in ethene C2H4 and ethyne C2H2 are described
a) Ethene C2H4 - Hybridisation take place for both carbon atoms in
ethene molecule is sp2 hybridisation
bull When both hybridised C is bonded together one of the hybridisedorbital overlapped directly between each other while the other two
hybridised orbitals overlapped directly with hydrogen atoms
z
atoms form a side-touch bond between each other and form anotherbond as shown in the diagram below
bull From the diagram with C=C there are two types of bond A sigma-
bond (σ-bond) is covalent bonds formed by orbitals overlapping end-
to-end with the electron density concentrated between the nuclei ofthe bonding atoms while the second type is called a pi bond (π-
bond) which is defined as a covalent bond formed by sideways
overlapping orbitals with electron density concentrated above and
below the plane of the nuclei of the bonding atoms
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital of 2s and two
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of atrigonal planar shape
Note that on unhybridised pz
orbital an electron is presence
Molecular shape Trigonal planar Angle between bond pair 1200
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b) Ethyne C H has the Lewis structure of H CequivC H which the
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b) Ethyne C2H2 has the Lewis structure of HminusCequivCminusH which the
bonding can be explain using sp hybridisation Table below
described how the hybridisation take place on each carbonatom and how the formation of triple bond occur
c) From the diagram formation of -CequivC- is due to the formation
of one sigma-bond by direct overlapping of one of the twohybrid orbitals on each C while the other two pi-bond bonds
are formed as a result of side-lapping of the each two
y z
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital from 2s and 2p
orbitals combined (hybrid) and
rearrange themselves to the
shape and orientation of a
linear shapeNote that on unhybridised py amp
pz orbitals electrons are
presence to form two
π-bonds Molecular shape Linear Angle between bond pair 1800
sp
pz py
5 Other examples of applications in valence bond theory includes the
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formation of nitrogen molecule N2 and hydrogen cyanide HCN
molecule
a) Nitrogen gas is Earths most abundant gas as it cover 78 of the
content of our air Nitrogen molecule is an inert gas thanks to its
short covalent bond and also its strong triple bond Therefore
a lot of heats are required to break the chemical bond of nitrogenbefore it can be applied in industries Using valence bond theory
the bonding of nitrogen molecule is explained in the diagram
b) So when the two hybridised nitrogen atom interacting amongeach other they hence formed a linear shape and two pi-bonds
(πminusbonds) are formed as a result of side-lapping of unhybridised
py and pz orbital respectively
N N
ExplanationEnergy level
diagramDiagram of orbitals
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diagram
a) Ground state
Both nitrogen whichact as the center atom
has the valence
electron of 2s22p3
b) Excited state One of the electron
from 2s is promote to 2p
orbitals - e ual ener
level
c) Hybridised state
One orbital of 2s and
one orbitals of 2p
combined (hybrid) and
rearrange themselves
to the shape and
orientation of a linear
shapeMolecular shape Linear
Angle between bond pair 1800
b) Whereas for hydrogen cyanide HCN both carbon atom and
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) y g y
nitrogen atom undergoes sp hybridisation in order to form a
linear structureExplanation
Energy level diagram for
carbon atom C
Energy level diagram for
nitrogen atom N
a) Ground state
Valence electronC 2s22p2
N 2s22p3
b) Excited state
One of the electron from
2s is promote to 2p
orbitals - equal energy
level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined
(hybrid) and form sp
hybridisation
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N
σσσσ
π
σσσσ
π
6 However there are a few limitation on valence bond theory such
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as when explaining the effect angle of bond-pair and bond-pair
electrons when there isare presence of lone pair electron in thecenter atom and also the difference of electronegativity
a) Bonding in ammonia NH3 and water H2O ~ Both NH3 and
H2O undergoes an arrangement similar to sp3 hybridisation
similar to that of C in methane Table below shows thehybridisation occur for nitrogen in ammonia and oxygen in water
Energy level diagram Energy level diagram
for nitrogen atom N for oxygen atom O
a) Ground state
Valence electron
N 2s22p3
O 2s2
2p4
b) Excited state
One of the electron
from 2s is promote to
2p orbitals
c) Hybridised state
One orbital of 2s and
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One orbital of 2s and
three orbitals from 2p
combined (hybrid) andrearrange themselves to
the shape and
orientation similar to that
of tetrahedral shape
Similar to arrangement in
tetrahedral
Similar to arrangement in
tetrahedral
of ammonia NH3 andwater H2O
Shape and angleShape trigonal pyramidal
Angle 1070
Shape bent
Angle 10450
Number of bond pair amp
lone pair electrons
Bond pair electron 3
Lone pair electron 1
Bond pair electron 2
Lone pair electron 2
bull The angle of bond pair - bond pair electrons in ammonia and0 0
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water are 1070 and 10450 respectively which is lesser than in
methane molecule (10950
) This can be explained by the factof the presence of lone-pair electrons in both ammonia and
water Since the lone pair - lone pair electron repulsion is
stronger than lone pair - bond pair electron repulsion than
bond pair - bond pair electron repulsion it is expected that therepulsion between lone pair - bond pair electrons in ammonia
is stronger than bond pair - bond pair electrons repulsion
hence caused the angle to squeeze to a smaller angle As
for water since there is a presence of lone pair - lone pair
electrons repulsion it results the bond pair - bond pair
electron repulsion to be much smaller hence caused the
angle to squeeze to a smaller angle Though valence bondtheory does not actually explain the hybridisation especially
when it relates to the repulsion occur involving lone pair
electron another theories shall be applied to study such
effect All these shall be discussed further in VSEPR theory
b) Effect of electronegativity and bonding angle between bond pair
- bond pair electrons ~
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- bond pair electrons
i As discussed earlier in the bonding of water (H2
O) the
molecular shape and angle is described in the diagram below
Sul hur which is also an element from Grou 16 formed h dro en
sulphide H2S when sulphur react with hydrogen However unlike
water the bonding angles is much smaller compare to water This is
due to the difference of electronegativity and also the bond length
between O and S in molecule Since O is more electronegative
compare to S the O-H bond is pulled closer toward O Furthermorethe bond length of O-H is shorter compare to S-H As a result the
bonding pair-bonding pair electrons repulse greater with each other in
H-O-H and caused the angle become greater
ii Another example is between NH3 and PH3 The orbital
f
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diagram for both NH3 and PH3 are described below Both
nitrogen N and phosphorous P are from the same groupwhich is Group 15
Similar to the above case P in phosphine is less
electronegative than N in ammonia and bond length of N-H
is shorter than P-H As a result H is pull closer to N and
repulsion between bonding of H-N-H is greater compare toH-P-H and caused the angle between H-N-H is greater
compare to H-P-H
324 Valence Shell Electron Pair Repulsion (VSEPR) Theory
1 M l l t i th th di i l t f t
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1 Molecular geometry is the three-dimensional arrangement of atoms
in a molecule A moleculersquos geometry affects its physical andchemical properties such as melting point boiling point density
and the types of reactions it undergoes The basic concept of
VSEPR theory is based on the three general rules below
a) As far as electron-pair repulsion is concerned double bonds andtriple bonds can be treated like single bonds This approximation is
good for qualitative purposes However you should realize that in
ldquo rdquo
because there are two or three bonds between two atoms theelectron density occupies more space
b) If a molecule has two or more resonance structures we can apply
the VSEPR model to any one of them Formal charges are usually
not shown
c) The order of repulsion strength of lone pair and bond pair are
lone-pair amp lone-pair electrons repulsion are the strongest
followed by lone-pair amp bond-pair electrons repulsion while
bond-pair amp bond-pair electrons repulsion is the weakest
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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2 There are also some stable covalent compounds which have
less than 8 electrons in their center atom (incomplete octet)
The center atoms are usually metals with great number ofvalence electrons with small atomic radius such as
beryllium boron and aluminium
BeCl2Be 2s2 Cl 3s23p5 BF3B 2 s22p1 F 2 s22p5 AlCl3Al 3s23p1 F 3 s23p5
These compound possessed stability due to their short bondlength between center atom with surround atom Furthermore
they can form resonance structure between the center atom
and surrounding atoms
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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3 There are also some molecules which contain an odd
number of electrons Among the most common compounds
are nitrogen monoxide (NO) and nitrogen dioxide (NO2)
Odd-electron molecules are sometimes called radicals
Nitrogen monoxide NO
N 2 s22p3 O 2 s22p4
Nitrogen dioxide NO2
N 2 s22p3 O 2 s22p4
any ra ca s are g y reac ve e reason s t at t ere
is a tendency for the unpaired electron to form a
covalent bond with an unpaired electron on another
molecule For example when two nitrogen dioxide
molecules collide they form dinitrogen tetroxide
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322 Dative bond
1 As shown in the Lewis structure of nitrate ion one of the N-Obond is drawn as rarr The bond rarr placed is called as dative
bond (also known as coordinative bond) where dative
bond is defined as a covalent bond in which one of the atoms
donate the lone pair electrons available Although theproperties of a coordinate covalent bond do not differ from
those of a normal covalent bond because all electrons are
a e no ma er w a e r source
2 Dative bond is usually applied for these few circumstances
below
a) To assist atom molecule ion that not yet achieved octet
configuration Making use of atom which has lone pair
electrons to those which are lack of electrons
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Sulphur dioxide SO2Sulphur trioxide SO3 Carbon monoxide CO
Ozone molecule O3 Water with hydrogen ion
S
OO
O
Ammonia with hydrogen ion Ammonia with boron trifluoride
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b) Formation of dimer - In order for some compounds which
have incomplete octet to achieve stability they tend to form
dimer or polymer among themselves by using dative bondTwo of the most common examples are aluminium trichloride
and beryllium dichloride
Monomer of AlCl3Dimer of aluminium chloride Al
2
Cl6
Monomer of BeCl2 Dimer of Be2Cl4 Polymer of (BeCl2)n
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c) Formation of coordination compounds - Coordination
compounds are substances that contain at least one complex
ion a species consisting of a central metal cation (either atransition metal or a main-group metal) that is bonded to
molecules andor anions called ligands via dative
(coordinative) bond
hexaaquacopper (II) ion [Cu(H2O)6]2+ tetraamminenickel (II) ion [Ni(NH3)4]
2+
Hexacyanoferrate (III) ion [Fe(CN)6]3- Trioxalatocobaltate(III)ion [Co(C2O4)3]
3-
3 2 3 Hybridisation Theory
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323 Hybridisation Theory
Valence bond theory
bull The basic principle of valence bond theory is that a covalentbond is formed when orbitals of two atoms overlap and the
overlapped region which is between the nuclei is occupied by
a pair of electrons The central themes of valence bond theory
derive from this principle
ndash Opposing spins of the electron pair ~ Stated in Paulis
orbitals has a maximum capacity of two electrons that musthave opposite spins For example when a covalent bond is
formed in molecule of hydrogen H2 the two 1s electrons of
two H atoms occupy the overlapping 1s orbitals and have
opposite spins
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bull Maximum overlap of bonding orbitals ~ The bond strength
depends on attraction of the nuclei for the shared electrons
so the greater the orbital overlap the stronger (morestable) the bond The extent of overlap depends on the
shapes and directions of the orbitals An s orbital is spherical
but p and d orbitals have more electron density in one
direction than in another Thus whenever possible a bond
involving p or d orbitals will be oriented in the direction that
maximizes overlap For example in hydrogen fluoride (HF)
bond the 1s orbital of H overlaps the half-filled 2p orbital of Falong the long axis of that orbital Any other direction would
result in less overlap and thus a weaker bond
H b idi ti f t i bit l T t f th
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bull Hybridisation of atomic orbitals ~ To account for the
bonding in simple diatomic molecules like HF we picture the
direct overlap of s and p orbitals of isolated atoms But howcan we account for the shapes of so many molecules and
polyatomic ions through the overlap of spherical s orbitals
dumbbell-shaped p orbitals and cloverleaf-shaped d orbitals
Linus Pauling proposed that the valence atomic orbitals in the
molecule are different from those in the isolated atoms The
spatial orientations of these new orbitals lead to more stable
bonds and are consistent with observed molecular shapesThe process of orbital mixing is called hybridisation and
the new atomic orbitals are called hybrid orbitals Two key
points about the number and type of hybrid orbitals are that
i The number of hybrid orbitals obtained equals the number ofatomic orbitals mixed
ii The type of hybrid orbitals obtained varies with the types of
atomic orbitals mixed
3 2 3 1 T f h b idi ti
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3231 Type of hybridisation
1 sp3 hybridisation ~ When four electron groups surround
the central atom the center atom involved must prepare fourorbitals with equal energies to overlap with the four
surrounding electron groups Valence Bond theory uses
hypothetical hybrid orbitals which are atomic orbitals
obtained when two or more non-equivalent orbitals of the
same atom combine in preparation for covalent bond
formation H bridisation is the term a lied to the mixin of
atomic orbitals in an atom (usually a central atom) togenerate a set of hybrid orbitals In the case of sp3 We can
generate four equivalent hybrid orbitals from the center
atom by mixing the s orbital and the three p orbitals
Explanation Energy level diagram Diagram of orbitals
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Explanation Energy level diagram Diagram of orbitals
a) Ground state
Carbon which act as thecenter atom has the valence
electron of 2s22p2
b) Excited state One of the electron from 2s
is promote to 2p orbitals -
equa energy eve
c) Hybridised state
One orbital of 2s and three
orbitals of 2p combined(hybrid) and rearrange
themselves to the shape
and orientation of a
tetrahedral shape
Molecular shape Tetrahedral
Angle between bond pair 10950
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2 sp2 hybridisation ~ When three electron groups surround
the central atom the center atom involved must prepare three
orbitals with equal energies to overlap with the three
surrounding electron groups In sp2 hybridisation three
equivalent (in terms of energy level) from the center atom by
mixing the one s orbital and the two p orbitals Using boron
trifluoride (BF3) as example sp2
hybridisation is explained
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3 sp hybridisation ~ When two electron groups surround
the central atom we observe a linear shape which
means that the bonding orbitals must have a linear
orientation VB theory explains this by proposing thatmixing two nonequivalent orbitals of a central atom one
s and one p gives rise to two equivalent sp hybrid
orbitals that lie 1800 apart
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Beryllium which act as the
center atom has the valence
electron of 2s2
b) Excited state
One of the electron from 2s ispromote to 2p orbitals - equal
energy level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of a
linear shape
Molecular shape Linear
Angle between bond pair 1800
4 The concept of hybridisation is also useful to explain molecules with
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p y p
doubletriple bonds By using the concept of the direct overlapping
orbitals and side-touch lapping orbitals the formation of multiple
bonds in ethene C2H4 and ethyne C2H2 are described
a) Ethene C2H4 - Hybridisation take place for both carbon atoms in
ethene molecule is sp2 hybridisation
bull When both hybridised C is bonded together one of the hybridisedorbital overlapped directly between each other while the other two
hybridised orbitals overlapped directly with hydrogen atoms
z
atoms form a side-touch bond between each other and form anotherbond as shown in the diagram below
bull From the diagram with C=C there are two types of bond A sigma-
bond (σ-bond) is covalent bonds formed by orbitals overlapping end-
to-end with the electron density concentrated between the nuclei ofthe bonding atoms while the second type is called a pi bond (π-
bond) which is defined as a covalent bond formed by sideways
overlapping orbitals with electron density concentrated above and
below the plane of the nuclei of the bonding atoms
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital of 2s and two
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of atrigonal planar shape
Note that on unhybridised pz
orbital an electron is presence
Molecular shape Trigonal planar Angle between bond pair 1200
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b) Ethyne C H has the Lewis structure of H CequivC H which the
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b) Ethyne C2H2 has the Lewis structure of HminusCequivCminusH which the
bonding can be explain using sp hybridisation Table below
described how the hybridisation take place on each carbonatom and how the formation of triple bond occur
c) From the diagram formation of -CequivC- is due to the formation
of one sigma-bond by direct overlapping of one of the twohybrid orbitals on each C while the other two pi-bond bonds
are formed as a result of side-lapping of the each two
y z
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital from 2s and 2p
orbitals combined (hybrid) and
rearrange themselves to the
shape and orientation of a
linear shapeNote that on unhybridised py amp
pz orbitals electrons are
presence to form two
π-bonds Molecular shape Linear Angle between bond pair 1800
sp
pz py
5 Other examples of applications in valence bond theory includes the
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formation of nitrogen molecule N2 and hydrogen cyanide HCN
molecule
a) Nitrogen gas is Earths most abundant gas as it cover 78 of the
content of our air Nitrogen molecule is an inert gas thanks to its
short covalent bond and also its strong triple bond Therefore
a lot of heats are required to break the chemical bond of nitrogenbefore it can be applied in industries Using valence bond theory
the bonding of nitrogen molecule is explained in the diagram
b) So when the two hybridised nitrogen atom interacting amongeach other they hence formed a linear shape and two pi-bonds
(πminusbonds) are formed as a result of side-lapping of unhybridised
py and pz orbital respectively
N N
ExplanationEnergy level
diagramDiagram of orbitals
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diagram
a) Ground state
Both nitrogen whichact as the center atom
has the valence
electron of 2s22p3
b) Excited state One of the electron
from 2s is promote to 2p
orbitals - e ual ener
level
c) Hybridised state
One orbital of 2s and
one orbitals of 2p
combined (hybrid) and
rearrange themselves
to the shape and
orientation of a linear
shapeMolecular shape Linear
Angle between bond pair 1800
b) Whereas for hydrogen cyanide HCN both carbon atom and
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) y g y
nitrogen atom undergoes sp hybridisation in order to form a
linear structureExplanation
Energy level diagram for
carbon atom C
Energy level diagram for
nitrogen atom N
a) Ground state
Valence electronC 2s22p2
N 2s22p3
b) Excited state
One of the electron from
2s is promote to 2p
orbitals - equal energy
level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined
(hybrid) and form sp
hybridisation
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N
σσσσ
π
σσσσ
π
6 However there are a few limitation on valence bond theory such
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as when explaining the effect angle of bond-pair and bond-pair
electrons when there isare presence of lone pair electron in thecenter atom and also the difference of electronegativity
a) Bonding in ammonia NH3 and water H2O ~ Both NH3 and
H2O undergoes an arrangement similar to sp3 hybridisation
similar to that of C in methane Table below shows thehybridisation occur for nitrogen in ammonia and oxygen in water
Energy level diagram Energy level diagram
for nitrogen atom N for oxygen atom O
a) Ground state
Valence electron
N 2s22p3
O 2s2
2p4
b) Excited state
One of the electron
from 2s is promote to
2p orbitals
c) Hybridised state
One orbital of 2s and
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One orbital of 2s and
three orbitals from 2p
combined (hybrid) andrearrange themselves to
the shape and
orientation similar to that
of tetrahedral shape
Similar to arrangement in
tetrahedral
Similar to arrangement in
tetrahedral
of ammonia NH3 andwater H2O
Shape and angleShape trigonal pyramidal
Angle 1070
Shape bent
Angle 10450
Number of bond pair amp
lone pair electrons
Bond pair electron 3
Lone pair electron 1
Bond pair electron 2
Lone pair electron 2
bull The angle of bond pair - bond pair electrons in ammonia and0 0
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water are 1070 and 10450 respectively which is lesser than in
methane molecule (10950
) This can be explained by the factof the presence of lone-pair electrons in both ammonia and
water Since the lone pair - lone pair electron repulsion is
stronger than lone pair - bond pair electron repulsion than
bond pair - bond pair electron repulsion it is expected that therepulsion between lone pair - bond pair electrons in ammonia
is stronger than bond pair - bond pair electrons repulsion
hence caused the angle to squeeze to a smaller angle As
for water since there is a presence of lone pair - lone pair
electrons repulsion it results the bond pair - bond pair
electron repulsion to be much smaller hence caused the
angle to squeeze to a smaller angle Though valence bondtheory does not actually explain the hybridisation especially
when it relates to the repulsion occur involving lone pair
electron another theories shall be applied to study such
effect All these shall be discussed further in VSEPR theory
b) Effect of electronegativity and bonding angle between bond pair
- bond pair electrons ~
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- bond pair electrons
i As discussed earlier in the bonding of water (H2
O) the
molecular shape and angle is described in the diagram below
Sul hur which is also an element from Grou 16 formed h dro en
sulphide H2S when sulphur react with hydrogen However unlike
water the bonding angles is much smaller compare to water This is
due to the difference of electronegativity and also the bond length
between O and S in molecule Since O is more electronegative
compare to S the O-H bond is pulled closer toward O Furthermorethe bond length of O-H is shorter compare to S-H As a result the
bonding pair-bonding pair electrons repulse greater with each other in
H-O-H and caused the angle become greater
ii Another example is between NH3 and PH3 The orbital
f
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diagram for both NH3 and PH3 are described below Both
nitrogen N and phosphorous P are from the same groupwhich is Group 15
Similar to the above case P in phosphine is less
electronegative than N in ammonia and bond length of N-H
is shorter than P-H As a result H is pull closer to N and
repulsion between bonding of H-N-H is greater compare toH-P-H and caused the angle between H-N-H is greater
compare to H-P-H
324 Valence Shell Electron Pair Repulsion (VSEPR) Theory
1 M l l t i th th di i l t f t
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1 Molecular geometry is the three-dimensional arrangement of atoms
in a molecule A moleculersquos geometry affects its physical andchemical properties such as melting point boiling point density
and the types of reactions it undergoes The basic concept of
VSEPR theory is based on the three general rules below
a) As far as electron-pair repulsion is concerned double bonds andtriple bonds can be treated like single bonds This approximation is
good for qualitative purposes However you should realize that in
ldquo rdquo
because there are two or three bonds between two atoms theelectron density occupies more space
b) If a molecule has two or more resonance structures we can apply
the VSEPR model to any one of them Formal charges are usually
not shown
c) The order of repulsion strength of lone pair and bond pair are
lone-pair amp lone-pair electrons repulsion are the strongest
followed by lone-pair amp bond-pair electrons repulsion while
bond-pair amp bond-pair electrons repulsion is the weakest
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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3 There are also some molecules which contain an odd
number of electrons Among the most common compounds
are nitrogen monoxide (NO) and nitrogen dioxide (NO2)
Odd-electron molecules are sometimes called radicals
Nitrogen monoxide NO
N 2 s22p3 O 2 s22p4
Nitrogen dioxide NO2
N 2 s22p3 O 2 s22p4
any ra ca s are g y reac ve e reason s t at t ere
is a tendency for the unpaired electron to form a
covalent bond with an unpaired electron on another
molecule For example when two nitrogen dioxide
molecules collide they form dinitrogen tetroxide
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322 Dative bond
1 As shown in the Lewis structure of nitrate ion one of the N-Obond is drawn as rarr The bond rarr placed is called as dative
bond (also known as coordinative bond) where dative
bond is defined as a covalent bond in which one of the atoms
donate the lone pair electrons available Although theproperties of a coordinate covalent bond do not differ from
those of a normal covalent bond because all electrons are
a e no ma er w a e r source
2 Dative bond is usually applied for these few circumstances
below
a) To assist atom molecule ion that not yet achieved octet
configuration Making use of atom which has lone pair
electrons to those which are lack of electrons
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Sulphur dioxide SO2Sulphur trioxide SO3 Carbon monoxide CO
Ozone molecule O3 Water with hydrogen ion
S
OO
O
Ammonia with hydrogen ion Ammonia with boron trifluoride
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b) Formation of dimer - In order for some compounds which
have incomplete octet to achieve stability they tend to form
dimer or polymer among themselves by using dative bondTwo of the most common examples are aluminium trichloride
and beryllium dichloride
Monomer of AlCl3Dimer of aluminium chloride Al
2
Cl6
Monomer of BeCl2 Dimer of Be2Cl4 Polymer of (BeCl2)n
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c) Formation of coordination compounds - Coordination
compounds are substances that contain at least one complex
ion a species consisting of a central metal cation (either atransition metal or a main-group metal) that is bonded to
molecules andor anions called ligands via dative
(coordinative) bond
hexaaquacopper (II) ion [Cu(H2O)6]2+ tetraamminenickel (II) ion [Ni(NH3)4]
2+
Hexacyanoferrate (III) ion [Fe(CN)6]3- Trioxalatocobaltate(III)ion [Co(C2O4)3]
3-
3 2 3 Hybridisation Theory
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323 Hybridisation Theory
Valence bond theory
bull The basic principle of valence bond theory is that a covalentbond is formed when orbitals of two atoms overlap and the
overlapped region which is between the nuclei is occupied by
a pair of electrons The central themes of valence bond theory
derive from this principle
ndash Opposing spins of the electron pair ~ Stated in Paulis
orbitals has a maximum capacity of two electrons that musthave opposite spins For example when a covalent bond is
formed in molecule of hydrogen H2 the two 1s electrons of
two H atoms occupy the overlapping 1s orbitals and have
opposite spins
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bull Maximum overlap of bonding orbitals ~ The bond strength
depends on attraction of the nuclei for the shared electrons
so the greater the orbital overlap the stronger (morestable) the bond The extent of overlap depends on the
shapes and directions of the orbitals An s orbital is spherical
but p and d orbitals have more electron density in one
direction than in another Thus whenever possible a bond
involving p or d orbitals will be oriented in the direction that
maximizes overlap For example in hydrogen fluoride (HF)
bond the 1s orbital of H overlaps the half-filled 2p orbital of Falong the long axis of that orbital Any other direction would
result in less overlap and thus a weaker bond
H b idi ti f t i bit l T t f th
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bull Hybridisation of atomic orbitals ~ To account for the
bonding in simple diatomic molecules like HF we picture the
direct overlap of s and p orbitals of isolated atoms But howcan we account for the shapes of so many molecules and
polyatomic ions through the overlap of spherical s orbitals
dumbbell-shaped p orbitals and cloverleaf-shaped d orbitals
Linus Pauling proposed that the valence atomic orbitals in the
molecule are different from those in the isolated atoms The
spatial orientations of these new orbitals lead to more stable
bonds and are consistent with observed molecular shapesThe process of orbital mixing is called hybridisation and
the new atomic orbitals are called hybrid orbitals Two key
points about the number and type of hybrid orbitals are that
i The number of hybrid orbitals obtained equals the number ofatomic orbitals mixed
ii The type of hybrid orbitals obtained varies with the types of
atomic orbitals mixed
3 2 3 1 T f h b idi ti
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3231 Type of hybridisation
1 sp3 hybridisation ~ When four electron groups surround
the central atom the center atom involved must prepare fourorbitals with equal energies to overlap with the four
surrounding electron groups Valence Bond theory uses
hypothetical hybrid orbitals which are atomic orbitals
obtained when two or more non-equivalent orbitals of the
same atom combine in preparation for covalent bond
formation H bridisation is the term a lied to the mixin of
atomic orbitals in an atom (usually a central atom) togenerate a set of hybrid orbitals In the case of sp3 We can
generate four equivalent hybrid orbitals from the center
atom by mixing the s orbital and the three p orbitals
Explanation Energy level diagram Diagram of orbitals
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Explanation Energy level diagram Diagram of orbitals
a) Ground state
Carbon which act as thecenter atom has the valence
electron of 2s22p2
b) Excited state One of the electron from 2s
is promote to 2p orbitals -
equa energy eve
c) Hybridised state
One orbital of 2s and three
orbitals of 2p combined(hybrid) and rearrange
themselves to the shape
and orientation of a
tetrahedral shape
Molecular shape Tetrahedral
Angle between bond pair 10950
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2 sp2 hybridisation ~ When three electron groups surround
the central atom the center atom involved must prepare three
orbitals with equal energies to overlap with the three
surrounding electron groups In sp2 hybridisation three
equivalent (in terms of energy level) from the center atom by
mixing the one s orbital and the two p orbitals Using boron
trifluoride (BF3) as example sp2
hybridisation is explained
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3 sp hybridisation ~ When two electron groups surround
the central atom we observe a linear shape which
means that the bonding orbitals must have a linear
orientation VB theory explains this by proposing thatmixing two nonequivalent orbitals of a central atom one
s and one p gives rise to two equivalent sp hybrid
orbitals that lie 1800 apart
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Beryllium which act as the
center atom has the valence
electron of 2s2
b) Excited state
One of the electron from 2s ispromote to 2p orbitals - equal
energy level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of a
linear shape
Molecular shape Linear
Angle between bond pair 1800
4 The concept of hybridisation is also useful to explain molecules with
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p y p
doubletriple bonds By using the concept of the direct overlapping
orbitals and side-touch lapping orbitals the formation of multiple
bonds in ethene C2H4 and ethyne C2H2 are described
a) Ethene C2H4 - Hybridisation take place for both carbon atoms in
ethene molecule is sp2 hybridisation
bull When both hybridised C is bonded together one of the hybridisedorbital overlapped directly between each other while the other two
hybridised orbitals overlapped directly with hydrogen atoms
z
atoms form a side-touch bond between each other and form anotherbond as shown in the diagram below
bull From the diagram with C=C there are two types of bond A sigma-
bond (σ-bond) is covalent bonds formed by orbitals overlapping end-
to-end with the electron density concentrated between the nuclei ofthe bonding atoms while the second type is called a pi bond (π-
bond) which is defined as a covalent bond formed by sideways
overlapping orbitals with electron density concentrated above and
below the plane of the nuclei of the bonding atoms
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital of 2s and two
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of atrigonal planar shape
Note that on unhybridised pz
orbital an electron is presence
Molecular shape Trigonal planar Angle between bond pair 1200
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b) Ethyne C H has the Lewis structure of H CequivC H which the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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b) Ethyne C2H2 has the Lewis structure of HminusCequivCminusH which the
bonding can be explain using sp hybridisation Table below
described how the hybridisation take place on each carbonatom and how the formation of triple bond occur
c) From the diagram formation of -CequivC- is due to the formation
of one sigma-bond by direct overlapping of one of the twohybrid orbitals on each C while the other two pi-bond bonds
are formed as a result of side-lapping of the each two
y z
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital from 2s and 2p
orbitals combined (hybrid) and
rearrange themselves to the
shape and orientation of a
linear shapeNote that on unhybridised py amp
pz orbitals electrons are
presence to form two
π-bonds Molecular shape Linear Angle between bond pair 1800
sp
pz py
5 Other examples of applications in valence bond theory includes the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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formation of nitrogen molecule N2 and hydrogen cyanide HCN
molecule
a) Nitrogen gas is Earths most abundant gas as it cover 78 of the
content of our air Nitrogen molecule is an inert gas thanks to its
short covalent bond and also its strong triple bond Therefore
a lot of heats are required to break the chemical bond of nitrogenbefore it can be applied in industries Using valence bond theory
the bonding of nitrogen molecule is explained in the diagram
b) So when the two hybridised nitrogen atom interacting amongeach other they hence formed a linear shape and two pi-bonds
(πminusbonds) are formed as a result of side-lapping of unhybridised
py and pz orbital respectively
N N
ExplanationEnergy level
diagramDiagram of orbitals
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diagram
a) Ground state
Both nitrogen whichact as the center atom
has the valence
electron of 2s22p3
b) Excited state One of the electron
from 2s is promote to 2p
orbitals - e ual ener
level
c) Hybridised state
One orbital of 2s and
one orbitals of 2p
combined (hybrid) and
rearrange themselves
to the shape and
orientation of a linear
shapeMolecular shape Linear
Angle between bond pair 1800
b) Whereas for hydrogen cyanide HCN both carbon atom and
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) y g y
nitrogen atom undergoes sp hybridisation in order to form a
linear structureExplanation
Energy level diagram for
carbon atom C
Energy level diagram for
nitrogen atom N
a) Ground state
Valence electronC 2s22p2
N 2s22p3
b) Excited state
One of the electron from
2s is promote to 2p
orbitals - equal energy
level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined
(hybrid) and form sp
hybridisation
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N
σσσσ
π
σσσσ
π
6 However there are a few limitation on valence bond theory such
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as when explaining the effect angle of bond-pair and bond-pair
electrons when there isare presence of lone pair electron in thecenter atom and also the difference of electronegativity
a) Bonding in ammonia NH3 and water H2O ~ Both NH3 and
H2O undergoes an arrangement similar to sp3 hybridisation
similar to that of C in methane Table below shows thehybridisation occur for nitrogen in ammonia and oxygen in water
Energy level diagram Energy level diagram
for nitrogen atom N for oxygen atom O
a) Ground state
Valence electron
N 2s22p3
O 2s2
2p4
b) Excited state
One of the electron
from 2s is promote to
2p orbitals
c) Hybridised state
One orbital of 2s and
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One orbital of 2s and
three orbitals from 2p
combined (hybrid) andrearrange themselves to
the shape and
orientation similar to that
of tetrahedral shape
Similar to arrangement in
tetrahedral
Similar to arrangement in
tetrahedral
of ammonia NH3 andwater H2O
Shape and angleShape trigonal pyramidal
Angle 1070
Shape bent
Angle 10450
Number of bond pair amp
lone pair electrons
Bond pair electron 3
Lone pair electron 1
Bond pair electron 2
Lone pair electron 2
bull The angle of bond pair - bond pair electrons in ammonia and0 0
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water are 1070 and 10450 respectively which is lesser than in
methane molecule (10950
) This can be explained by the factof the presence of lone-pair electrons in both ammonia and
water Since the lone pair - lone pair electron repulsion is
stronger than lone pair - bond pair electron repulsion than
bond pair - bond pair electron repulsion it is expected that therepulsion between lone pair - bond pair electrons in ammonia
is stronger than bond pair - bond pair electrons repulsion
hence caused the angle to squeeze to a smaller angle As
for water since there is a presence of lone pair - lone pair
electrons repulsion it results the bond pair - bond pair
electron repulsion to be much smaller hence caused the
angle to squeeze to a smaller angle Though valence bondtheory does not actually explain the hybridisation especially
when it relates to the repulsion occur involving lone pair
electron another theories shall be applied to study such
effect All these shall be discussed further in VSEPR theory
b) Effect of electronegativity and bonding angle between bond pair
- bond pair electrons ~
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- bond pair electrons
i As discussed earlier in the bonding of water (H2
O) the
molecular shape and angle is described in the diagram below
Sul hur which is also an element from Grou 16 formed h dro en
sulphide H2S when sulphur react with hydrogen However unlike
water the bonding angles is much smaller compare to water This is
due to the difference of electronegativity and also the bond length
between O and S in molecule Since O is more electronegative
compare to S the O-H bond is pulled closer toward O Furthermorethe bond length of O-H is shorter compare to S-H As a result the
bonding pair-bonding pair electrons repulse greater with each other in
H-O-H and caused the angle become greater
ii Another example is between NH3 and PH3 The orbital
f
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diagram for both NH3 and PH3 are described below Both
nitrogen N and phosphorous P are from the same groupwhich is Group 15
Similar to the above case P in phosphine is less
electronegative than N in ammonia and bond length of N-H
is shorter than P-H As a result H is pull closer to N and
repulsion between bonding of H-N-H is greater compare toH-P-H and caused the angle between H-N-H is greater
compare to H-P-H
324 Valence Shell Electron Pair Repulsion (VSEPR) Theory
1 M l l t i th th di i l t f t
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1 Molecular geometry is the three-dimensional arrangement of atoms
in a molecule A moleculersquos geometry affects its physical andchemical properties such as melting point boiling point density
and the types of reactions it undergoes The basic concept of
VSEPR theory is based on the three general rules below
a) As far as electron-pair repulsion is concerned double bonds andtriple bonds can be treated like single bonds This approximation is
good for qualitative purposes However you should realize that in
ldquo rdquo
because there are two or three bonds between two atoms theelectron density occupies more space
b) If a molecule has two or more resonance structures we can apply
the VSEPR model to any one of them Formal charges are usually
not shown
c) The order of repulsion strength of lone pair and bond pair are
lone-pair amp lone-pair electrons repulsion are the strongest
followed by lone-pair amp bond-pair electrons repulsion while
bond-pair amp bond-pair electrons repulsion is the weakest
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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322 Dative bond
1 As shown in the Lewis structure of nitrate ion one of the N-Obond is drawn as rarr The bond rarr placed is called as dative
bond (also known as coordinative bond) where dative
bond is defined as a covalent bond in which one of the atoms
donate the lone pair electrons available Although theproperties of a coordinate covalent bond do not differ from
those of a normal covalent bond because all electrons are
a e no ma er w a e r source
2 Dative bond is usually applied for these few circumstances
below
a) To assist atom molecule ion that not yet achieved octet
configuration Making use of atom which has lone pair
electrons to those which are lack of electrons
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Sulphur dioxide SO2Sulphur trioxide SO3 Carbon monoxide CO
Ozone molecule O3 Water with hydrogen ion
S
OO
O
Ammonia with hydrogen ion Ammonia with boron trifluoride
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b) Formation of dimer - In order for some compounds which
have incomplete octet to achieve stability they tend to form
dimer or polymer among themselves by using dative bondTwo of the most common examples are aluminium trichloride
and beryllium dichloride
Monomer of AlCl3Dimer of aluminium chloride Al
2
Cl6
Monomer of BeCl2 Dimer of Be2Cl4 Polymer of (BeCl2)n
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c) Formation of coordination compounds - Coordination
compounds are substances that contain at least one complex
ion a species consisting of a central metal cation (either atransition metal or a main-group metal) that is bonded to
molecules andor anions called ligands via dative
(coordinative) bond
hexaaquacopper (II) ion [Cu(H2O)6]2+ tetraamminenickel (II) ion [Ni(NH3)4]
2+
Hexacyanoferrate (III) ion [Fe(CN)6]3- Trioxalatocobaltate(III)ion [Co(C2O4)3]
3-
3 2 3 Hybridisation Theory
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323 Hybridisation Theory
Valence bond theory
bull The basic principle of valence bond theory is that a covalentbond is formed when orbitals of two atoms overlap and the
overlapped region which is between the nuclei is occupied by
a pair of electrons The central themes of valence bond theory
derive from this principle
ndash Opposing spins of the electron pair ~ Stated in Paulis
orbitals has a maximum capacity of two electrons that musthave opposite spins For example when a covalent bond is
formed in molecule of hydrogen H2 the two 1s electrons of
two H atoms occupy the overlapping 1s orbitals and have
opposite spins
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bull Maximum overlap of bonding orbitals ~ The bond strength
depends on attraction of the nuclei for the shared electrons
so the greater the orbital overlap the stronger (morestable) the bond The extent of overlap depends on the
shapes and directions of the orbitals An s orbital is spherical
but p and d orbitals have more electron density in one
direction than in another Thus whenever possible a bond
involving p or d orbitals will be oriented in the direction that
maximizes overlap For example in hydrogen fluoride (HF)
bond the 1s orbital of H overlaps the half-filled 2p orbital of Falong the long axis of that orbital Any other direction would
result in less overlap and thus a weaker bond
H b idi ti f t i bit l T t f th
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bull Hybridisation of atomic orbitals ~ To account for the
bonding in simple diatomic molecules like HF we picture the
direct overlap of s and p orbitals of isolated atoms But howcan we account for the shapes of so many molecules and
polyatomic ions through the overlap of spherical s orbitals
dumbbell-shaped p orbitals and cloverleaf-shaped d orbitals
Linus Pauling proposed that the valence atomic orbitals in the
molecule are different from those in the isolated atoms The
spatial orientations of these new orbitals lead to more stable
bonds and are consistent with observed molecular shapesThe process of orbital mixing is called hybridisation and
the new atomic orbitals are called hybrid orbitals Two key
points about the number and type of hybrid orbitals are that
i The number of hybrid orbitals obtained equals the number ofatomic orbitals mixed
ii The type of hybrid orbitals obtained varies with the types of
atomic orbitals mixed
3 2 3 1 T f h b idi ti
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3231 Type of hybridisation
1 sp3 hybridisation ~ When four electron groups surround
the central atom the center atom involved must prepare fourorbitals with equal energies to overlap with the four
surrounding electron groups Valence Bond theory uses
hypothetical hybrid orbitals which are atomic orbitals
obtained when two or more non-equivalent orbitals of the
same atom combine in preparation for covalent bond
formation H bridisation is the term a lied to the mixin of
atomic orbitals in an atom (usually a central atom) togenerate a set of hybrid orbitals In the case of sp3 We can
generate four equivalent hybrid orbitals from the center
atom by mixing the s orbital and the three p orbitals
Explanation Energy level diagram Diagram of orbitals
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Explanation Energy level diagram Diagram of orbitals
a) Ground state
Carbon which act as thecenter atom has the valence
electron of 2s22p2
b) Excited state One of the electron from 2s
is promote to 2p orbitals -
equa energy eve
c) Hybridised state
One orbital of 2s and three
orbitals of 2p combined(hybrid) and rearrange
themselves to the shape
and orientation of a
tetrahedral shape
Molecular shape Tetrahedral
Angle between bond pair 10950
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2 sp2 hybridisation ~ When three electron groups surround
the central atom the center atom involved must prepare three
orbitals with equal energies to overlap with the three
surrounding electron groups In sp2 hybridisation three
equivalent (in terms of energy level) from the center atom by
mixing the one s orbital and the two p orbitals Using boron
trifluoride (BF3) as example sp2
hybridisation is explained
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3 sp hybridisation ~ When two electron groups surround
the central atom we observe a linear shape which
means that the bonding orbitals must have a linear
orientation VB theory explains this by proposing thatmixing two nonequivalent orbitals of a central atom one
s and one p gives rise to two equivalent sp hybrid
orbitals that lie 1800 apart
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Beryllium which act as the
center atom has the valence
electron of 2s2
b) Excited state
One of the electron from 2s ispromote to 2p orbitals - equal
energy level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of a
linear shape
Molecular shape Linear
Angle between bond pair 1800
4 The concept of hybridisation is also useful to explain molecules with
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p y p
doubletriple bonds By using the concept of the direct overlapping
orbitals and side-touch lapping orbitals the formation of multiple
bonds in ethene C2H4 and ethyne C2H2 are described
a) Ethene C2H4 - Hybridisation take place for both carbon atoms in
ethene molecule is sp2 hybridisation
bull When both hybridised C is bonded together one of the hybridisedorbital overlapped directly between each other while the other two
hybridised orbitals overlapped directly with hydrogen atoms
z
atoms form a side-touch bond between each other and form anotherbond as shown in the diagram below
bull From the diagram with C=C there are two types of bond A sigma-
bond (σ-bond) is covalent bonds formed by orbitals overlapping end-
to-end with the electron density concentrated between the nuclei ofthe bonding atoms while the second type is called a pi bond (π-
bond) which is defined as a covalent bond formed by sideways
overlapping orbitals with electron density concentrated above and
below the plane of the nuclei of the bonding atoms
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital of 2s and two
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of atrigonal planar shape
Note that on unhybridised pz
orbital an electron is presence
Molecular shape Trigonal planar Angle between bond pair 1200
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b) Ethyne C H has the Lewis structure of H CequivC H which the
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b) Ethyne C2H2 has the Lewis structure of HminusCequivCminusH which the
bonding can be explain using sp hybridisation Table below
described how the hybridisation take place on each carbonatom and how the formation of triple bond occur
c) From the diagram formation of -CequivC- is due to the formation
of one sigma-bond by direct overlapping of one of the twohybrid orbitals on each C while the other two pi-bond bonds
are formed as a result of side-lapping of the each two
y z
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital from 2s and 2p
orbitals combined (hybrid) and
rearrange themselves to the
shape and orientation of a
linear shapeNote that on unhybridised py amp
pz orbitals electrons are
presence to form two
π-bonds Molecular shape Linear Angle between bond pair 1800
sp
pz py
5 Other examples of applications in valence bond theory includes the
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formation of nitrogen molecule N2 and hydrogen cyanide HCN
molecule
a) Nitrogen gas is Earths most abundant gas as it cover 78 of the
content of our air Nitrogen molecule is an inert gas thanks to its
short covalent bond and also its strong triple bond Therefore
a lot of heats are required to break the chemical bond of nitrogenbefore it can be applied in industries Using valence bond theory
the bonding of nitrogen molecule is explained in the diagram
b) So when the two hybridised nitrogen atom interacting amongeach other they hence formed a linear shape and two pi-bonds
(πminusbonds) are formed as a result of side-lapping of unhybridised
py and pz orbital respectively
N N
ExplanationEnergy level
diagramDiagram of orbitals
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diagram
a) Ground state
Both nitrogen whichact as the center atom
has the valence
electron of 2s22p3
b) Excited state One of the electron
from 2s is promote to 2p
orbitals - e ual ener
level
c) Hybridised state
One orbital of 2s and
one orbitals of 2p
combined (hybrid) and
rearrange themselves
to the shape and
orientation of a linear
shapeMolecular shape Linear
Angle between bond pair 1800
b) Whereas for hydrogen cyanide HCN both carbon atom and
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) y g y
nitrogen atom undergoes sp hybridisation in order to form a
linear structureExplanation
Energy level diagram for
carbon atom C
Energy level diagram for
nitrogen atom N
a) Ground state
Valence electronC 2s22p2
N 2s22p3
b) Excited state
One of the electron from
2s is promote to 2p
orbitals - equal energy
level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined
(hybrid) and form sp
hybridisation
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N
σσσσ
π
σσσσ
π
6 However there are a few limitation on valence bond theory such
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as when explaining the effect angle of bond-pair and bond-pair
electrons when there isare presence of lone pair electron in thecenter atom and also the difference of electronegativity
a) Bonding in ammonia NH3 and water H2O ~ Both NH3 and
H2O undergoes an arrangement similar to sp3 hybridisation
similar to that of C in methane Table below shows thehybridisation occur for nitrogen in ammonia and oxygen in water
Energy level diagram Energy level diagram
for nitrogen atom N for oxygen atom O
a) Ground state
Valence electron
N 2s22p3
O 2s2
2p4
b) Excited state
One of the electron
from 2s is promote to
2p orbitals
c) Hybridised state
One orbital of 2s and
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One orbital of 2s and
three orbitals from 2p
combined (hybrid) andrearrange themselves to
the shape and
orientation similar to that
of tetrahedral shape
Similar to arrangement in
tetrahedral
Similar to arrangement in
tetrahedral
of ammonia NH3 andwater H2O
Shape and angleShape trigonal pyramidal
Angle 1070
Shape bent
Angle 10450
Number of bond pair amp
lone pair electrons
Bond pair electron 3
Lone pair electron 1
Bond pair electron 2
Lone pair electron 2
bull The angle of bond pair - bond pair electrons in ammonia and0 0
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water are 1070 and 10450 respectively which is lesser than in
methane molecule (10950
) This can be explained by the factof the presence of lone-pair electrons in both ammonia and
water Since the lone pair - lone pair electron repulsion is
stronger than lone pair - bond pair electron repulsion than
bond pair - bond pair electron repulsion it is expected that therepulsion between lone pair - bond pair electrons in ammonia
is stronger than bond pair - bond pair electrons repulsion
hence caused the angle to squeeze to a smaller angle As
for water since there is a presence of lone pair - lone pair
electrons repulsion it results the bond pair - bond pair
electron repulsion to be much smaller hence caused the
angle to squeeze to a smaller angle Though valence bondtheory does not actually explain the hybridisation especially
when it relates to the repulsion occur involving lone pair
electron another theories shall be applied to study such
effect All these shall be discussed further in VSEPR theory
b) Effect of electronegativity and bonding angle between bond pair
- bond pair electrons ~
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- bond pair electrons
i As discussed earlier in the bonding of water (H2
O) the
molecular shape and angle is described in the diagram below
Sul hur which is also an element from Grou 16 formed h dro en
sulphide H2S when sulphur react with hydrogen However unlike
water the bonding angles is much smaller compare to water This is
due to the difference of electronegativity and also the bond length
between O and S in molecule Since O is more electronegative
compare to S the O-H bond is pulled closer toward O Furthermorethe bond length of O-H is shorter compare to S-H As a result the
bonding pair-bonding pair electrons repulse greater with each other in
H-O-H and caused the angle become greater
ii Another example is between NH3 and PH3 The orbital
f
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diagram for both NH3 and PH3 are described below Both
nitrogen N and phosphorous P are from the same groupwhich is Group 15
Similar to the above case P in phosphine is less
electronegative than N in ammonia and bond length of N-H
is shorter than P-H As a result H is pull closer to N and
repulsion between bonding of H-N-H is greater compare toH-P-H and caused the angle between H-N-H is greater
compare to H-P-H
324 Valence Shell Electron Pair Repulsion (VSEPR) Theory
1 M l l t i th th di i l t f t
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1 Molecular geometry is the three-dimensional arrangement of atoms
in a molecule A moleculersquos geometry affects its physical andchemical properties such as melting point boiling point density
and the types of reactions it undergoes The basic concept of
VSEPR theory is based on the three general rules below
a) As far as electron-pair repulsion is concerned double bonds andtriple bonds can be treated like single bonds This approximation is
good for qualitative purposes However you should realize that in
ldquo rdquo
because there are two or three bonds between two atoms theelectron density occupies more space
b) If a molecule has two or more resonance structures we can apply
the VSEPR model to any one of them Formal charges are usually
not shown
c) The order of repulsion strength of lone pair and bond pair are
lone-pair amp lone-pair electrons repulsion are the strongest
followed by lone-pair amp bond-pair electrons repulsion while
bond-pair amp bond-pair electrons repulsion is the weakest
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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Sulphur dioxide SO2Sulphur trioxide SO3 Carbon monoxide CO
Ozone molecule O3 Water with hydrogen ion
S
OO
O
Ammonia with hydrogen ion Ammonia with boron trifluoride
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b) Formation of dimer - In order for some compounds which
have incomplete octet to achieve stability they tend to form
dimer or polymer among themselves by using dative bondTwo of the most common examples are aluminium trichloride
and beryllium dichloride
Monomer of AlCl3Dimer of aluminium chloride Al
2
Cl6
Monomer of BeCl2 Dimer of Be2Cl4 Polymer of (BeCl2)n
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c) Formation of coordination compounds - Coordination
compounds are substances that contain at least one complex
ion a species consisting of a central metal cation (either atransition metal or a main-group metal) that is bonded to
molecules andor anions called ligands via dative
(coordinative) bond
hexaaquacopper (II) ion [Cu(H2O)6]2+ tetraamminenickel (II) ion [Ni(NH3)4]
2+
Hexacyanoferrate (III) ion [Fe(CN)6]3- Trioxalatocobaltate(III)ion [Co(C2O4)3]
3-
3 2 3 Hybridisation Theory
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323 Hybridisation Theory
Valence bond theory
bull The basic principle of valence bond theory is that a covalentbond is formed when orbitals of two atoms overlap and the
overlapped region which is between the nuclei is occupied by
a pair of electrons The central themes of valence bond theory
derive from this principle
ndash Opposing spins of the electron pair ~ Stated in Paulis
orbitals has a maximum capacity of two electrons that musthave opposite spins For example when a covalent bond is
formed in molecule of hydrogen H2 the two 1s electrons of
two H atoms occupy the overlapping 1s orbitals and have
opposite spins
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bull Maximum overlap of bonding orbitals ~ The bond strength
depends on attraction of the nuclei for the shared electrons
so the greater the orbital overlap the stronger (morestable) the bond The extent of overlap depends on the
shapes and directions of the orbitals An s orbital is spherical
but p and d orbitals have more electron density in one
direction than in another Thus whenever possible a bond
involving p or d orbitals will be oriented in the direction that
maximizes overlap For example in hydrogen fluoride (HF)
bond the 1s orbital of H overlaps the half-filled 2p orbital of Falong the long axis of that orbital Any other direction would
result in less overlap and thus a weaker bond
H b idi ti f t i bit l T t f th
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bull Hybridisation of atomic orbitals ~ To account for the
bonding in simple diatomic molecules like HF we picture the
direct overlap of s and p orbitals of isolated atoms But howcan we account for the shapes of so many molecules and
polyatomic ions through the overlap of spherical s orbitals
dumbbell-shaped p orbitals and cloverleaf-shaped d orbitals
Linus Pauling proposed that the valence atomic orbitals in the
molecule are different from those in the isolated atoms The
spatial orientations of these new orbitals lead to more stable
bonds and are consistent with observed molecular shapesThe process of orbital mixing is called hybridisation and
the new atomic orbitals are called hybrid orbitals Two key
points about the number and type of hybrid orbitals are that
i The number of hybrid orbitals obtained equals the number ofatomic orbitals mixed
ii The type of hybrid orbitals obtained varies with the types of
atomic orbitals mixed
3 2 3 1 T f h b idi ti
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3231 Type of hybridisation
1 sp3 hybridisation ~ When four electron groups surround
the central atom the center atom involved must prepare fourorbitals with equal energies to overlap with the four
surrounding electron groups Valence Bond theory uses
hypothetical hybrid orbitals which are atomic orbitals
obtained when two or more non-equivalent orbitals of the
same atom combine in preparation for covalent bond
formation H bridisation is the term a lied to the mixin of
atomic orbitals in an atom (usually a central atom) togenerate a set of hybrid orbitals In the case of sp3 We can
generate four equivalent hybrid orbitals from the center
atom by mixing the s orbital and the three p orbitals
Explanation Energy level diagram Diagram of orbitals
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Explanation Energy level diagram Diagram of orbitals
a) Ground state
Carbon which act as thecenter atom has the valence
electron of 2s22p2
b) Excited state One of the electron from 2s
is promote to 2p orbitals -
equa energy eve
c) Hybridised state
One orbital of 2s and three
orbitals of 2p combined(hybrid) and rearrange
themselves to the shape
and orientation of a
tetrahedral shape
Molecular shape Tetrahedral
Angle between bond pair 10950
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2 sp2 hybridisation ~ When three electron groups surround
the central atom the center atom involved must prepare three
orbitals with equal energies to overlap with the three
surrounding electron groups In sp2 hybridisation three
equivalent (in terms of energy level) from the center atom by
mixing the one s orbital and the two p orbitals Using boron
trifluoride (BF3) as example sp2
hybridisation is explained
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8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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3 sp hybridisation ~ When two electron groups surround
the central atom we observe a linear shape which
means that the bonding orbitals must have a linear
orientation VB theory explains this by proposing thatmixing two nonequivalent orbitals of a central atom one
s and one p gives rise to two equivalent sp hybrid
orbitals that lie 1800 apart
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Beryllium which act as the
center atom has the valence
electron of 2s2
b) Excited state
One of the electron from 2s ispromote to 2p orbitals - equal
energy level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of a
linear shape
Molecular shape Linear
Angle between bond pair 1800
4 The concept of hybridisation is also useful to explain molecules with
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p y p
doubletriple bonds By using the concept of the direct overlapping
orbitals and side-touch lapping orbitals the formation of multiple
bonds in ethene C2H4 and ethyne C2H2 are described
a) Ethene C2H4 - Hybridisation take place for both carbon atoms in
ethene molecule is sp2 hybridisation
bull When both hybridised C is bonded together one of the hybridisedorbital overlapped directly between each other while the other two
hybridised orbitals overlapped directly with hydrogen atoms
z
atoms form a side-touch bond between each other and form anotherbond as shown in the diagram below
bull From the diagram with C=C there are two types of bond A sigma-
bond (σ-bond) is covalent bonds formed by orbitals overlapping end-
to-end with the electron density concentrated between the nuclei ofthe bonding atoms while the second type is called a pi bond (π-
bond) which is defined as a covalent bond formed by sideways
overlapping orbitals with electron density concentrated above and
below the plane of the nuclei of the bonding atoms
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital of 2s and two
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of atrigonal planar shape
Note that on unhybridised pz
orbital an electron is presence
Molecular shape Trigonal planar Angle between bond pair 1200
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b) Ethyne C H has the Lewis structure of H CequivC H which the
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b) Ethyne C2H2 has the Lewis structure of HminusCequivCminusH which the
bonding can be explain using sp hybridisation Table below
described how the hybridisation take place on each carbonatom and how the formation of triple bond occur
c) From the diagram formation of -CequivC- is due to the formation
of one sigma-bond by direct overlapping of one of the twohybrid orbitals on each C while the other two pi-bond bonds
are formed as a result of side-lapping of the each two
y z
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital from 2s and 2p
orbitals combined (hybrid) and
rearrange themselves to the
shape and orientation of a
linear shapeNote that on unhybridised py amp
pz orbitals electrons are
presence to form two
π-bonds Molecular shape Linear Angle between bond pair 1800
sp
pz py
5 Other examples of applications in valence bond theory includes the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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formation of nitrogen molecule N2 and hydrogen cyanide HCN
molecule
a) Nitrogen gas is Earths most abundant gas as it cover 78 of the
content of our air Nitrogen molecule is an inert gas thanks to its
short covalent bond and also its strong triple bond Therefore
a lot of heats are required to break the chemical bond of nitrogenbefore it can be applied in industries Using valence bond theory
the bonding of nitrogen molecule is explained in the diagram
b) So when the two hybridised nitrogen atom interacting amongeach other they hence formed a linear shape and two pi-bonds
(πminusbonds) are formed as a result of side-lapping of unhybridised
py and pz orbital respectively
N N
ExplanationEnergy level
diagramDiagram of orbitals
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diagram
a) Ground state
Both nitrogen whichact as the center atom
has the valence
electron of 2s22p3
b) Excited state One of the electron
from 2s is promote to 2p
orbitals - e ual ener
level
c) Hybridised state
One orbital of 2s and
one orbitals of 2p
combined (hybrid) and
rearrange themselves
to the shape and
orientation of a linear
shapeMolecular shape Linear
Angle between bond pair 1800
b) Whereas for hydrogen cyanide HCN both carbon atom and
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) y g y
nitrogen atom undergoes sp hybridisation in order to form a
linear structureExplanation
Energy level diagram for
carbon atom C
Energy level diagram for
nitrogen atom N
a) Ground state
Valence electronC 2s22p2
N 2s22p3
b) Excited state
One of the electron from
2s is promote to 2p
orbitals - equal energy
level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined
(hybrid) and form sp
hybridisation
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N
σσσσ
π
σσσσ
π
6 However there are a few limitation on valence bond theory such
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as when explaining the effect angle of bond-pair and bond-pair
electrons when there isare presence of lone pair electron in thecenter atom and also the difference of electronegativity
a) Bonding in ammonia NH3 and water H2O ~ Both NH3 and
H2O undergoes an arrangement similar to sp3 hybridisation
similar to that of C in methane Table below shows thehybridisation occur for nitrogen in ammonia and oxygen in water
Energy level diagram Energy level diagram
for nitrogen atom N for oxygen atom O
a) Ground state
Valence electron
N 2s22p3
O 2s2
2p4
b) Excited state
One of the electron
from 2s is promote to
2p orbitals
c) Hybridised state
One orbital of 2s and
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One orbital of 2s and
three orbitals from 2p
combined (hybrid) andrearrange themselves to
the shape and
orientation similar to that
of tetrahedral shape
Similar to arrangement in
tetrahedral
Similar to arrangement in
tetrahedral
of ammonia NH3 andwater H2O
Shape and angleShape trigonal pyramidal
Angle 1070
Shape bent
Angle 10450
Number of bond pair amp
lone pair electrons
Bond pair electron 3
Lone pair electron 1
Bond pair electron 2
Lone pair electron 2
bull The angle of bond pair - bond pair electrons in ammonia and0 0
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water are 1070 and 10450 respectively which is lesser than in
methane molecule (10950
) This can be explained by the factof the presence of lone-pair electrons in both ammonia and
water Since the lone pair - lone pair electron repulsion is
stronger than lone pair - bond pair electron repulsion than
bond pair - bond pair electron repulsion it is expected that therepulsion between lone pair - bond pair electrons in ammonia
is stronger than bond pair - bond pair electrons repulsion
hence caused the angle to squeeze to a smaller angle As
for water since there is a presence of lone pair - lone pair
electrons repulsion it results the bond pair - bond pair
electron repulsion to be much smaller hence caused the
angle to squeeze to a smaller angle Though valence bondtheory does not actually explain the hybridisation especially
when it relates to the repulsion occur involving lone pair
electron another theories shall be applied to study such
effect All these shall be discussed further in VSEPR theory
b) Effect of electronegativity and bonding angle between bond pair
- bond pair electrons ~
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- bond pair electrons
i As discussed earlier in the bonding of water (H2
O) the
molecular shape and angle is described in the diagram below
Sul hur which is also an element from Grou 16 formed h dro en
sulphide H2S when sulphur react with hydrogen However unlike
water the bonding angles is much smaller compare to water This is
due to the difference of electronegativity and also the bond length
between O and S in molecule Since O is more electronegative
compare to S the O-H bond is pulled closer toward O Furthermorethe bond length of O-H is shorter compare to S-H As a result the
bonding pair-bonding pair electrons repulse greater with each other in
H-O-H and caused the angle become greater
ii Another example is between NH3 and PH3 The orbital
f
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diagram for both NH3 and PH3 are described below Both
nitrogen N and phosphorous P are from the same groupwhich is Group 15
Similar to the above case P in phosphine is less
electronegative than N in ammonia and bond length of N-H
is shorter than P-H As a result H is pull closer to N and
repulsion between bonding of H-N-H is greater compare toH-P-H and caused the angle between H-N-H is greater
compare to H-P-H
324 Valence Shell Electron Pair Repulsion (VSEPR) Theory
1 M l l t i th th di i l t f t
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1 Molecular geometry is the three-dimensional arrangement of atoms
in a molecule A moleculersquos geometry affects its physical andchemical properties such as melting point boiling point density
and the types of reactions it undergoes The basic concept of
VSEPR theory is based on the three general rules below
a) As far as electron-pair repulsion is concerned double bonds andtriple bonds can be treated like single bonds This approximation is
good for qualitative purposes However you should realize that in
ldquo rdquo
because there are two or three bonds between two atoms theelectron density occupies more space
b) If a molecule has two or more resonance structures we can apply
the VSEPR model to any one of them Formal charges are usually
not shown
c) The order of repulsion strength of lone pair and bond pair are
lone-pair amp lone-pair electrons repulsion are the strongest
followed by lone-pair amp bond-pair electrons repulsion while
bond-pair amp bond-pair electrons repulsion is the weakest
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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b) Formation of dimer - In order for some compounds which
have incomplete octet to achieve stability they tend to form
dimer or polymer among themselves by using dative bondTwo of the most common examples are aluminium trichloride
and beryllium dichloride
Monomer of AlCl3Dimer of aluminium chloride Al
2
Cl6
Monomer of BeCl2 Dimer of Be2Cl4 Polymer of (BeCl2)n
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c) Formation of coordination compounds - Coordination
compounds are substances that contain at least one complex
ion a species consisting of a central metal cation (either atransition metal or a main-group metal) that is bonded to
molecules andor anions called ligands via dative
(coordinative) bond
hexaaquacopper (II) ion [Cu(H2O)6]2+ tetraamminenickel (II) ion [Ni(NH3)4]
2+
Hexacyanoferrate (III) ion [Fe(CN)6]3- Trioxalatocobaltate(III)ion [Co(C2O4)3]
3-
3 2 3 Hybridisation Theory
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323 Hybridisation Theory
Valence bond theory
bull The basic principle of valence bond theory is that a covalentbond is formed when orbitals of two atoms overlap and the
overlapped region which is between the nuclei is occupied by
a pair of electrons The central themes of valence bond theory
derive from this principle
ndash Opposing spins of the electron pair ~ Stated in Paulis
orbitals has a maximum capacity of two electrons that musthave opposite spins For example when a covalent bond is
formed in molecule of hydrogen H2 the two 1s electrons of
two H atoms occupy the overlapping 1s orbitals and have
opposite spins
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bull Maximum overlap of bonding orbitals ~ The bond strength
depends on attraction of the nuclei for the shared electrons
so the greater the orbital overlap the stronger (morestable) the bond The extent of overlap depends on the
shapes and directions of the orbitals An s orbital is spherical
but p and d orbitals have more electron density in one
direction than in another Thus whenever possible a bond
involving p or d orbitals will be oriented in the direction that
maximizes overlap For example in hydrogen fluoride (HF)
bond the 1s orbital of H overlaps the half-filled 2p orbital of Falong the long axis of that orbital Any other direction would
result in less overlap and thus a weaker bond
H b idi ti f t i bit l T t f th
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bull Hybridisation of atomic orbitals ~ To account for the
bonding in simple diatomic molecules like HF we picture the
direct overlap of s and p orbitals of isolated atoms But howcan we account for the shapes of so many molecules and
polyatomic ions through the overlap of spherical s orbitals
dumbbell-shaped p orbitals and cloverleaf-shaped d orbitals
Linus Pauling proposed that the valence atomic orbitals in the
molecule are different from those in the isolated atoms The
spatial orientations of these new orbitals lead to more stable
bonds and are consistent with observed molecular shapesThe process of orbital mixing is called hybridisation and
the new atomic orbitals are called hybrid orbitals Two key
points about the number and type of hybrid orbitals are that
i The number of hybrid orbitals obtained equals the number ofatomic orbitals mixed
ii The type of hybrid orbitals obtained varies with the types of
atomic orbitals mixed
3 2 3 1 T f h b idi ti
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3231 Type of hybridisation
1 sp3 hybridisation ~ When four electron groups surround
the central atom the center atom involved must prepare fourorbitals with equal energies to overlap with the four
surrounding electron groups Valence Bond theory uses
hypothetical hybrid orbitals which are atomic orbitals
obtained when two or more non-equivalent orbitals of the
same atom combine in preparation for covalent bond
formation H bridisation is the term a lied to the mixin of
atomic orbitals in an atom (usually a central atom) togenerate a set of hybrid orbitals In the case of sp3 We can
generate four equivalent hybrid orbitals from the center
atom by mixing the s orbital and the three p orbitals
Explanation Energy level diagram Diagram of orbitals
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Explanation Energy level diagram Diagram of orbitals
a) Ground state
Carbon which act as thecenter atom has the valence
electron of 2s22p2
b) Excited state One of the electron from 2s
is promote to 2p orbitals -
equa energy eve
c) Hybridised state
One orbital of 2s and three
orbitals of 2p combined(hybrid) and rearrange
themselves to the shape
and orientation of a
tetrahedral shape
Molecular shape Tetrahedral
Angle between bond pair 10950
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2 sp2 hybridisation ~ When three electron groups surround
the central atom the center atom involved must prepare three
orbitals with equal energies to overlap with the three
surrounding electron groups In sp2 hybridisation three
equivalent (in terms of energy level) from the center atom by
mixing the one s orbital and the two p orbitals Using boron
trifluoride (BF3) as example sp2
hybridisation is explained
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3 sp hybridisation ~ When two electron groups surround
the central atom we observe a linear shape which
means that the bonding orbitals must have a linear
orientation VB theory explains this by proposing thatmixing two nonequivalent orbitals of a central atom one
s and one p gives rise to two equivalent sp hybrid
orbitals that lie 1800 apart
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Beryllium which act as the
center atom has the valence
electron of 2s2
b) Excited state
One of the electron from 2s ispromote to 2p orbitals - equal
energy level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of a
linear shape
Molecular shape Linear
Angle between bond pair 1800
4 The concept of hybridisation is also useful to explain molecules with
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p y p
doubletriple bonds By using the concept of the direct overlapping
orbitals and side-touch lapping orbitals the formation of multiple
bonds in ethene C2H4 and ethyne C2H2 are described
a) Ethene C2H4 - Hybridisation take place for both carbon atoms in
ethene molecule is sp2 hybridisation
bull When both hybridised C is bonded together one of the hybridisedorbital overlapped directly between each other while the other two
hybridised orbitals overlapped directly with hydrogen atoms
z
atoms form a side-touch bond between each other and form anotherbond as shown in the diagram below
bull From the diagram with C=C there are two types of bond A sigma-
bond (σ-bond) is covalent bonds formed by orbitals overlapping end-
to-end with the electron density concentrated between the nuclei ofthe bonding atoms while the second type is called a pi bond (π-
bond) which is defined as a covalent bond formed by sideways
overlapping orbitals with electron density concentrated above and
below the plane of the nuclei of the bonding atoms
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital of 2s and two
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of atrigonal planar shape
Note that on unhybridised pz
orbital an electron is presence
Molecular shape Trigonal planar Angle between bond pair 1200
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b) Ethyne C H has the Lewis structure of H CequivC H which the
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b) Ethyne C2H2 has the Lewis structure of HminusCequivCminusH which the
bonding can be explain using sp hybridisation Table below
described how the hybridisation take place on each carbonatom and how the formation of triple bond occur
c) From the diagram formation of -CequivC- is due to the formation
of one sigma-bond by direct overlapping of one of the twohybrid orbitals on each C while the other two pi-bond bonds
are formed as a result of side-lapping of the each two
y z
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital from 2s and 2p
orbitals combined (hybrid) and
rearrange themselves to the
shape and orientation of a
linear shapeNote that on unhybridised py amp
pz orbitals electrons are
presence to form two
π-bonds Molecular shape Linear Angle between bond pair 1800
sp
pz py
5 Other examples of applications in valence bond theory includes the
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formation of nitrogen molecule N2 and hydrogen cyanide HCN
molecule
a) Nitrogen gas is Earths most abundant gas as it cover 78 of the
content of our air Nitrogen molecule is an inert gas thanks to its
short covalent bond and also its strong triple bond Therefore
a lot of heats are required to break the chemical bond of nitrogenbefore it can be applied in industries Using valence bond theory
the bonding of nitrogen molecule is explained in the diagram
b) So when the two hybridised nitrogen atom interacting amongeach other they hence formed a linear shape and two pi-bonds
(πminusbonds) are formed as a result of side-lapping of unhybridised
py and pz orbital respectively
N N
ExplanationEnergy level
diagramDiagram of orbitals
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diagram
a) Ground state
Both nitrogen whichact as the center atom
has the valence
electron of 2s22p3
b) Excited state One of the electron
from 2s is promote to 2p
orbitals - e ual ener
level
c) Hybridised state
One orbital of 2s and
one orbitals of 2p
combined (hybrid) and
rearrange themselves
to the shape and
orientation of a linear
shapeMolecular shape Linear
Angle between bond pair 1800
b) Whereas for hydrogen cyanide HCN both carbon atom and
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) y g y
nitrogen atom undergoes sp hybridisation in order to form a
linear structureExplanation
Energy level diagram for
carbon atom C
Energy level diagram for
nitrogen atom N
a) Ground state
Valence electronC 2s22p2
N 2s22p3
b) Excited state
One of the electron from
2s is promote to 2p
orbitals - equal energy
level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined
(hybrid) and form sp
hybridisation
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N
σσσσ
π
σσσσ
π
6 However there are a few limitation on valence bond theory such
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as when explaining the effect angle of bond-pair and bond-pair
electrons when there isare presence of lone pair electron in thecenter atom and also the difference of electronegativity
a) Bonding in ammonia NH3 and water H2O ~ Both NH3 and
H2O undergoes an arrangement similar to sp3 hybridisation
similar to that of C in methane Table below shows thehybridisation occur for nitrogen in ammonia and oxygen in water
Energy level diagram Energy level diagram
for nitrogen atom N for oxygen atom O
a) Ground state
Valence electron
N 2s22p3
O 2s2
2p4
b) Excited state
One of the electron
from 2s is promote to
2p orbitals
c) Hybridised state
One orbital of 2s and
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One orbital of 2s and
three orbitals from 2p
combined (hybrid) andrearrange themselves to
the shape and
orientation similar to that
of tetrahedral shape
Similar to arrangement in
tetrahedral
Similar to arrangement in
tetrahedral
of ammonia NH3 andwater H2O
Shape and angleShape trigonal pyramidal
Angle 1070
Shape bent
Angle 10450
Number of bond pair amp
lone pair electrons
Bond pair electron 3
Lone pair electron 1
Bond pair electron 2
Lone pair electron 2
bull The angle of bond pair - bond pair electrons in ammonia and0 0
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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water are 1070 and 10450 respectively which is lesser than in
methane molecule (10950
) This can be explained by the factof the presence of lone-pair electrons in both ammonia and
water Since the lone pair - lone pair electron repulsion is
stronger than lone pair - bond pair electron repulsion than
bond pair - bond pair electron repulsion it is expected that therepulsion between lone pair - bond pair electrons in ammonia
is stronger than bond pair - bond pair electrons repulsion
hence caused the angle to squeeze to a smaller angle As
for water since there is a presence of lone pair - lone pair
electrons repulsion it results the bond pair - bond pair
electron repulsion to be much smaller hence caused the
angle to squeeze to a smaller angle Though valence bondtheory does not actually explain the hybridisation especially
when it relates to the repulsion occur involving lone pair
electron another theories shall be applied to study such
effect All these shall be discussed further in VSEPR theory
b) Effect of electronegativity and bonding angle between bond pair
- bond pair electrons ~
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- bond pair electrons
i As discussed earlier in the bonding of water (H2
O) the
molecular shape and angle is described in the diagram below
Sul hur which is also an element from Grou 16 formed h dro en
sulphide H2S when sulphur react with hydrogen However unlike
water the bonding angles is much smaller compare to water This is
due to the difference of electronegativity and also the bond length
between O and S in molecule Since O is more electronegative
compare to S the O-H bond is pulled closer toward O Furthermorethe bond length of O-H is shorter compare to S-H As a result the
bonding pair-bonding pair electrons repulse greater with each other in
H-O-H and caused the angle become greater
ii Another example is between NH3 and PH3 The orbital
f
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diagram for both NH3 and PH3 are described below Both
nitrogen N and phosphorous P are from the same groupwhich is Group 15
Similar to the above case P in phosphine is less
electronegative than N in ammonia and bond length of N-H
is shorter than P-H As a result H is pull closer to N and
repulsion between bonding of H-N-H is greater compare toH-P-H and caused the angle between H-N-H is greater
compare to H-P-H
324 Valence Shell Electron Pair Repulsion (VSEPR) Theory
1 M l l t i th th di i l t f t
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1 Molecular geometry is the three-dimensional arrangement of atoms
in a molecule A moleculersquos geometry affects its physical andchemical properties such as melting point boiling point density
and the types of reactions it undergoes The basic concept of
VSEPR theory is based on the three general rules below
a) As far as electron-pair repulsion is concerned double bonds andtriple bonds can be treated like single bonds This approximation is
good for qualitative purposes However you should realize that in
ldquo rdquo
because there are two or three bonds between two atoms theelectron density occupies more space
b) If a molecule has two or more resonance structures we can apply
the VSEPR model to any one of them Formal charges are usually
not shown
c) The order of repulsion strength of lone pair and bond pair are
lone-pair amp lone-pair electrons repulsion are the strongest
followed by lone-pair amp bond-pair electrons repulsion while
bond-pair amp bond-pair electrons repulsion is the weakest
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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c) Formation of coordination compounds - Coordination
compounds are substances that contain at least one complex
ion a species consisting of a central metal cation (either atransition metal or a main-group metal) that is bonded to
molecules andor anions called ligands via dative
(coordinative) bond
hexaaquacopper (II) ion [Cu(H2O)6]2+ tetraamminenickel (II) ion [Ni(NH3)4]
2+
Hexacyanoferrate (III) ion [Fe(CN)6]3- Trioxalatocobaltate(III)ion [Co(C2O4)3]
3-
3 2 3 Hybridisation Theory
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323 Hybridisation Theory
Valence bond theory
bull The basic principle of valence bond theory is that a covalentbond is formed when orbitals of two atoms overlap and the
overlapped region which is between the nuclei is occupied by
a pair of electrons The central themes of valence bond theory
derive from this principle
ndash Opposing spins of the electron pair ~ Stated in Paulis
orbitals has a maximum capacity of two electrons that musthave opposite spins For example when a covalent bond is
formed in molecule of hydrogen H2 the two 1s electrons of
two H atoms occupy the overlapping 1s orbitals and have
opposite spins
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bull Maximum overlap of bonding orbitals ~ The bond strength
depends on attraction of the nuclei for the shared electrons
so the greater the orbital overlap the stronger (morestable) the bond The extent of overlap depends on the
shapes and directions of the orbitals An s orbital is spherical
but p and d orbitals have more electron density in one
direction than in another Thus whenever possible a bond
involving p or d orbitals will be oriented in the direction that
maximizes overlap For example in hydrogen fluoride (HF)
bond the 1s orbital of H overlaps the half-filled 2p orbital of Falong the long axis of that orbital Any other direction would
result in less overlap and thus a weaker bond
H b idi ti f t i bit l T t f th
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bull Hybridisation of atomic orbitals ~ To account for the
bonding in simple diatomic molecules like HF we picture the
direct overlap of s and p orbitals of isolated atoms But howcan we account for the shapes of so many molecules and
polyatomic ions through the overlap of spherical s orbitals
dumbbell-shaped p orbitals and cloverleaf-shaped d orbitals
Linus Pauling proposed that the valence atomic orbitals in the
molecule are different from those in the isolated atoms The
spatial orientations of these new orbitals lead to more stable
bonds and are consistent with observed molecular shapesThe process of orbital mixing is called hybridisation and
the new atomic orbitals are called hybrid orbitals Two key
points about the number and type of hybrid orbitals are that
i The number of hybrid orbitals obtained equals the number ofatomic orbitals mixed
ii The type of hybrid orbitals obtained varies with the types of
atomic orbitals mixed
3 2 3 1 T f h b idi ti
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3231 Type of hybridisation
1 sp3 hybridisation ~ When four electron groups surround
the central atom the center atom involved must prepare fourorbitals with equal energies to overlap with the four
surrounding electron groups Valence Bond theory uses
hypothetical hybrid orbitals which are atomic orbitals
obtained when two or more non-equivalent orbitals of the
same atom combine in preparation for covalent bond
formation H bridisation is the term a lied to the mixin of
atomic orbitals in an atom (usually a central atom) togenerate a set of hybrid orbitals In the case of sp3 We can
generate four equivalent hybrid orbitals from the center
atom by mixing the s orbital and the three p orbitals
Explanation Energy level diagram Diagram of orbitals
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Explanation Energy level diagram Diagram of orbitals
a) Ground state
Carbon which act as thecenter atom has the valence
electron of 2s22p2
b) Excited state One of the electron from 2s
is promote to 2p orbitals -
equa energy eve
c) Hybridised state
One orbital of 2s and three
orbitals of 2p combined(hybrid) and rearrange
themselves to the shape
and orientation of a
tetrahedral shape
Molecular shape Tetrahedral
Angle between bond pair 10950
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2 sp2 hybridisation ~ When three electron groups surround
the central atom the center atom involved must prepare three
orbitals with equal energies to overlap with the three
surrounding electron groups In sp2 hybridisation three
equivalent (in terms of energy level) from the center atom by
mixing the one s orbital and the two p orbitals Using boron
trifluoride (BF3) as example sp2
hybridisation is explained
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8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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3 sp hybridisation ~ When two electron groups surround
the central atom we observe a linear shape which
means that the bonding orbitals must have a linear
orientation VB theory explains this by proposing thatmixing two nonequivalent orbitals of a central atom one
s and one p gives rise to two equivalent sp hybrid
orbitals that lie 1800 apart
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Beryllium which act as the
center atom has the valence
electron of 2s2
b) Excited state
One of the electron from 2s ispromote to 2p orbitals - equal
energy level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of a
linear shape
Molecular shape Linear
Angle between bond pair 1800
4 The concept of hybridisation is also useful to explain molecules with
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p y p
doubletriple bonds By using the concept of the direct overlapping
orbitals and side-touch lapping orbitals the formation of multiple
bonds in ethene C2H4 and ethyne C2H2 are described
a) Ethene C2H4 - Hybridisation take place for both carbon atoms in
ethene molecule is sp2 hybridisation
bull When both hybridised C is bonded together one of the hybridisedorbital overlapped directly between each other while the other two
hybridised orbitals overlapped directly with hydrogen atoms
z
atoms form a side-touch bond between each other and form anotherbond as shown in the diagram below
bull From the diagram with C=C there are two types of bond A sigma-
bond (σ-bond) is covalent bonds formed by orbitals overlapping end-
to-end with the electron density concentrated between the nuclei ofthe bonding atoms while the second type is called a pi bond (π-
bond) which is defined as a covalent bond formed by sideways
overlapping orbitals with electron density concentrated above and
below the plane of the nuclei of the bonding atoms
Explanation Energy level diagram Diagram of orbitals
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital of 2s and two
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of atrigonal planar shape
Note that on unhybridised pz
orbital an electron is presence
Molecular shape Trigonal planar Angle between bond pair 1200
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b) Ethyne C H has the Lewis structure of H CequivC H which the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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b) Ethyne C2H2 has the Lewis structure of HminusCequivCminusH which the
bonding can be explain using sp hybridisation Table below
described how the hybridisation take place on each carbonatom and how the formation of triple bond occur
c) From the diagram formation of -CequivC- is due to the formation
of one sigma-bond by direct overlapping of one of the twohybrid orbitals on each C while the other two pi-bond bonds
are formed as a result of side-lapping of the each two
y z
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital from 2s and 2p
orbitals combined (hybrid) and
rearrange themselves to the
shape and orientation of a
linear shapeNote that on unhybridised py amp
pz orbitals electrons are
presence to form two
π-bonds Molecular shape Linear Angle between bond pair 1800
sp
pz py
5 Other examples of applications in valence bond theory includes the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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formation of nitrogen molecule N2 and hydrogen cyanide HCN
molecule
a) Nitrogen gas is Earths most abundant gas as it cover 78 of the
content of our air Nitrogen molecule is an inert gas thanks to its
short covalent bond and also its strong triple bond Therefore
a lot of heats are required to break the chemical bond of nitrogenbefore it can be applied in industries Using valence bond theory
the bonding of nitrogen molecule is explained in the diagram
b) So when the two hybridised nitrogen atom interacting amongeach other they hence formed a linear shape and two pi-bonds
(πminusbonds) are formed as a result of side-lapping of unhybridised
py and pz orbital respectively
N N
ExplanationEnergy level
diagramDiagram of orbitals
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diagram
a) Ground state
Both nitrogen whichact as the center atom
has the valence
electron of 2s22p3
b) Excited state One of the electron
from 2s is promote to 2p
orbitals - e ual ener
level
c) Hybridised state
One orbital of 2s and
one orbitals of 2p
combined (hybrid) and
rearrange themselves
to the shape and
orientation of a linear
shapeMolecular shape Linear
Angle between bond pair 1800
b) Whereas for hydrogen cyanide HCN both carbon atom and
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) y g y
nitrogen atom undergoes sp hybridisation in order to form a
linear structureExplanation
Energy level diagram for
carbon atom C
Energy level diagram for
nitrogen atom N
a) Ground state
Valence electronC 2s22p2
N 2s22p3
b) Excited state
One of the electron from
2s is promote to 2p
orbitals - equal energy
level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined
(hybrid) and form sp
hybridisation
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N
σσσσ
π
σσσσ
π
6 However there are a few limitation on valence bond theory such
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as when explaining the effect angle of bond-pair and bond-pair
electrons when there isare presence of lone pair electron in thecenter atom and also the difference of electronegativity
a) Bonding in ammonia NH3 and water H2O ~ Both NH3 and
H2O undergoes an arrangement similar to sp3 hybridisation
similar to that of C in methane Table below shows thehybridisation occur for nitrogen in ammonia and oxygen in water
Energy level diagram Energy level diagram
for nitrogen atom N for oxygen atom O
a) Ground state
Valence electron
N 2s22p3
O 2s2
2p4
b) Excited state
One of the electron
from 2s is promote to
2p orbitals
c) Hybridised state
One orbital of 2s and
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One orbital of 2s and
three orbitals from 2p
combined (hybrid) andrearrange themselves to
the shape and
orientation similar to that
of tetrahedral shape
Similar to arrangement in
tetrahedral
Similar to arrangement in
tetrahedral
of ammonia NH3 andwater H2O
Shape and angleShape trigonal pyramidal
Angle 1070
Shape bent
Angle 10450
Number of bond pair amp
lone pair electrons
Bond pair electron 3
Lone pair electron 1
Bond pair electron 2
Lone pair electron 2
bull The angle of bond pair - bond pair electrons in ammonia and0 0
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water are 1070 and 10450 respectively which is lesser than in
methane molecule (10950
) This can be explained by the factof the presence of lone-pair electrons in both ammonia and
water Since the lone pair - lone pair electron repulsion is
stronger than lone pair - bond pair electron repulsion than
bond pair - bond pair electron repulsion it is expected that therepulsion between lone pair - bond pair electrons in ammonia
is stronger than bond pair - bond pair electrons repulsion
hence caused the angle to squeeze to a smaller angle As
for water since there is a presence of lone pair - lone pair
electrons repulsion it results the bond pair - bond pair
electron repulsion to be much smaller hence caused the
angle to squeeze to a smaller angle Though valence bondtheory does not actually explain the hybridisation especially
when it relates to the repulsion occur involving lone pair
electron another theories shall be applied to study such
effect All these shall be discussed further in VSEPR theory
b) Effect of electronegativity and bonding angle between bond pair
- bond pair electrons ~
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- bond pair electrons
i As discussed earlier in the bonding of water (H2
O) the
molecular shape and angle is described in the diagram below
Sul hur which is also an element from Grou 16 formed h dro en
sulphide H2S when sulphur react with hydrogen However unlike
water the bonding angles is much smaller compare to water This is
due to the difference of electronegativity and also the bond length
between O and S in molecule Since O is more electronegative
compare to S the O-H bond is pulled closer toward O Furthermorethe bond length of O-H is shorter compare to S-H As a result the
bonding pair-bonding pair electrons repulse greater with each other in
H-O-H and caused the angle become greater
ii Another example is between NH3 and PH3 The orbital
f
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diagram for both NH3 and PH3 are described below Both
nitrogen N and phosphorous P are from the same groupwhich is Group 15
Similar to the above case P in phosphine is less
electronegative than N in ammonia and bond length of N-H
is shorter than P-H As a result H is pull closer to N and
repulsion between bonding of H-N-H is greater compare toH-P-H and caused the angle between H-N-H is greater
compare to H-P-H
324 Valence Shell Electron Pair Repulsion (VSEPR) Theory
1 M l l t i th th di i l t f t
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1 Molecular geometry is the three-dimensional arrangement of atoms
in a molecule A moleculersquos geometry affects its physical andchemical properties such as melting point boiling point density
and the types of reactions it undergoes The basic concept of
VSEPR theory is based on the three general rules below
a) As far as electron-pair repulsion is concerned double bonds andtriple bonds can be treated like single bonds This approximation is
good for qualitative purposes However you should realize that in
ldquo rdquo
because there are two or three bonds between two atoms theelectron density occupies more space
b) If a molecule has two or more resonance structures we can apply
the VSEPR model to any one of them Formal charges are usually
not shown
c) The order of repulsion strength of lone pair and bond pair are
lone-pair amp lone-pair electrons repulsion are the strongest
followed by lone-pair amp bond-pair electrons repulsion while
bond-pair amp bond-pair electrons repulsion is the weakest
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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323 Hybridisation Theory
Valence bond theory
bull The basic principle of valence bond theory is that a covalentbond is formed when orbitals of two atoms overlap and the
overlapped region which is between the nuclei is occupied by
a pair of electrons The central themes of valence bond theory
derive from this principle
ndash Opposing spins of the electron pair ~ Stated in Paulis
orbitals has a maximum capacity of two electrons that musthave opposite spins For example when a covalent bond is
formed in molecule of hydrogen H2 the two 1s electrons of
two H atoms occupy the overlapping 1s orbitals and have
opposite spins
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bull Maximum overlap of bonding orbitals ~ The bond strength
depends on attraction of the nuclei for the shared electrons
so the greater the orbital overlap the stronger (morestable) the bond The extent of overlap depends on the
shapes and directions of the orbitals An s orbital is spherical
but p and d orbitals have more electron density in one
direction than in another Thus whenever possible a bond
involving p or d orbitals will be oriented in the direction that
maximizes overlap For example in hydrogen fluoride (HF)
bond the 1s orbital of H overlaps the half-filled 2p orbital of Falong the long axis of that orbital Any other direction would
result in less overlap and thus a weaker bond
H b idi ti f t i bit l T t f th
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bull Hybridisation of atomic orbitals ~ To account for the
bonding in simple diatomic molecules like HF we picture the
direct overlap of s and p orbitals of isolated atoms But howcan we account for the shapes of so many molecules and
polyatomic ions through the overlap of spherical s orbitals
dumbbell-shaped p orbitals and cloverleaf-shaped d orbitals
Linus Pauling proposed that the valence atomic orbitals in the
molecule are different from those in the isolated atoms The
spatial orientations of these new orbitals lead to more stable
bonds and are consistent with observed molecular shapesThe process of orbital mixing is called hybridisation and
the new atomic orbitals are called hybrid orbitals Two key
points about the number and type of hybrid orbitals are that
i The number of hybrid orbitals obtained equals the number ofatomic orbitals mixed
ii The type of hybrid orbitals obtained varies with the types of
atomic orbitals mixed
3 2 3 1 T f h b idi ti
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3231 Type of hybridisation
1 sp3 hybridisation ~ When four electron groups surround
the central atom the center atom involved must prepare fourorbitals with equal energies to overlap with the four
surrounding electron groups Valence Bond theory uses
hypothetical hybrid orbitals which are atomic orbitals
obtained when two or more non-equivalent orbitals of the
same atom combine in preparation for covalent bond
formation H bridisation is the term a lied to the mixin of
atomic orbitals in an atom (usually a central atom) togenerate a set of hybrid orbitals In the case of sp3 We can
generate four equivalent hybrid orbitals from the center
atom by mixing the s orbital and the three p orbitals
Explanation Energy level diagram Diagram of orbitals
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Explanation Energy level diagram Diagram of orbitals
a) Ground state
Carbon which act as thecenter atom has the valence
electron of 2s22p2
b) Excited state One of the electron from 2s
is promote to 2p orbitals -
equa energy eve
c) Hybridised state
One orbital of 2s and three
orbitals of 2p combined(hybrid) and rearrange
themselves to the shape
and orientation of a
tetrahedral shape
Molecular shape Tetrahedral
Angle between bond pair 10950
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2 sp2 hybridisation ~ When three electron groups surround
the central atom the center atom involved must prepare three
orbitals with equal energies to overlap with the three
surrounding electron groups In sp2 hybridisation three
equivalent (in terms of energy level) from the center atom by
mixing the one s orbital and the two p orbitals Using boron
trifluoride (BF3) as example sp2
hybridisation is explained
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3 sp hybridisation ~ When two electron groups surround
the central atom we observe a linear shape which
means that the bonding orbitals must have a linear
orientation VB theory explains this by proposing thatmixing two nonequivalent orbitals of a central atom one
s and one p gives rise to two equivalent sp hybrid
orbitals that lie 1800 apart
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Beryllium which act as the
center atom has the valence
electron of 2s2
b) Excited state
One of the electron from 2s ispromote to 2p orbitals - equal
energy level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of a
linear shape
Molecular shape Linear
Angle between bond pair 1800
4 The concept of hybridisation is also useful to explain molecules with
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p y p
doubletriple bonds By using the concept of the direct overlapping
orbitals and side-touch lapping orbitals the formation of multiple
bonds in ethene C2H4 and ethyne C2H2 are described
a) Ethene C2H4 - Hybridisation take place for both carbon atoms in
ethene molecule is sp2 hybridisation
bull When both hybridised C is bonded together one of the hybridisedorbital overlapped directly between each other while the other two
hybridised orbitals overlapped directly with hydrogen atoms
z
atoms form a side-touch bond between each other and form anotherbond as shown in the diagram below
bull From the diagram with C=C there are two types of bond A sigma-
bond (σ-bond) is covalent bonds formed by orbitals overlapping end-
to-end with the electron density concentrated between the nuclei ofthe bonding atoms while the second type is called a pi bond (π-
bond) which is defined as a covalent bond formed by sideways
overlapping orbitals with electron density concentrated above and
below the plane of the nuclei of the bonding atoms
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital of 2s and two
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of atrigonal planar shape
Note that on unhybridised pz
orbital an electron is presence
Molecular shape Trigonal planar Angle between bond pair 1200
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b) Ethyne C H has the Lewis structure of H CequivC H which the
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b) Ethyne C2H2 has the Lewis structure of HminusCequivCminusH which the
bonding can be explain using sp hybridisation Table below
described how the hybridisation take place on each carbonatom and how the formation of triple bond occur
c) From the diagram formation of -CequivC- is due to the formation
of one sigma-bond by direct overlapping of one of the twohybrid orbitals on each C while the other two pi-bond bonds
are formed as a result of side-lapping of the each two
y z
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital from 2s and 2p
orbitals combined (hybrid) and
rearrange themselves to the
shape and orientation of a
linear shapeNote that on unhybridised py amp
pz orbitals electrons are
presence to form two
π-bonds Molecular shape Linear Angle between bond pair 1800
sp
pz py
5 Other examples of applications in valence bond theory includes the
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formation of nitrogen molecule N2 and hydrogen cyanide HCN
molecule
a) Nitrogen gas is Earths most abundant gas as it cover 78 of the
content of our air Nitrogen molecule is an inert gas thanks to its
short covalent bond and also its strong triple bond Therefore
a lot of heats are required to break the chemical bond of nitrogenbefore it can be applied in industries Using valence bond theory
the bonding of nitrogen molecule is explained in the diagram
b) So when the two hybridised nitrogen atom interacting amongeach other they hence formed a linear shape and two pi-bonds
(πminusbonds) are formed as a result of side-lapping of unhybridised
py and pz orbital respectively
N N
ExplanationEnergy level
diagramDiagram of orbitals
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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diagram
a) Ground state
Both nitrogen whichact as the center atom
has the valence
electron of 2s22p3
b) Excited state One of the electron
from 2s is promote to 2p
orbitals - e ual ener
level
c) Hybridised state
One orbital of 2s and
one orbitals of 2p
combined (hybrid) and
rearrange themselves
to the shape and
orientation of a linear
shapeMolecular shape Linear
Angle between bond pair 1800
b) Whereas for hydrogen cyanide HCN both carbon atom and
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) y g y
nitrogen atom undergoes sp hybridisation in order to form a
linear structureExplanation
Energy level diagram for
carbon atom C
Energy level diagram for
nitrogen atom N
a) Ground state
Valence electronC 2s22p2
N 2s22p3
b) Excited state
One of the electron from
2s is promote to 2p
orbitals - equal energy
level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined
(hybrid) and form sp
hybridisation
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N
σσσσ
π
σσσσ
π
6 However there are a few limitation on valence bond theory such
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as when explaining the effect angle of bond-pair and bond-pair
electrons when there isare presence of lone pair electron in thecenter atom and also the difference of electronegativity
a) Bonding in ammonia NH3 and water H2O ~ Both NH3 and
H2O undergoes an arrangement similar to sp3 hybridisation
similar to that of C in methane Table below shows thehybridisation occur for nitrogen in ammonia and oxygen in water
Energy level diagram Energy level diagram
for nitrogen atom N for oxygen atom O
a) Ground state
Valence electron
N 2s22p3
O 2s2
2p4
b) Excited state
One of the electron
from 2s is promote to
2p orbitals
c) Hybridised state
One orbital of 2s and
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One orbital of 2s and
three orbitals from 2p
combined (hybrid) andrearrange themselves to
the shape and
orientation similar to that
of tetrahedral shape
Similar to arrangement in
tetrahedral
Similar to arrangement in
tetrahedral
of ammonia NH3 andwater H2O
Shape and angleShape trigonal pyramidal
Angle 1070
Shape bent
Angle 10450
Number of bond pair amp
lone pair electrons
Bond pair electron 3
Lone pair electron 1
Bond pair electron 2
Lone pair electron 2
bull The angle of bond pair - bond pair electrons in ammonia and0 0
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water are 1070 and 10450 respectively which is lesser than in
methane molecule (10950
) This can be explained by the factof the presence of lone-pair electrons in both ammonia and
water Since the lone pair - lone pair electron repulsion is
stronger than lone pair - bond pair electron repulsion than
bond pair - bond pair electron repulsion it is expected that therepulsion between lone pair - bond pair electrons in ammonia
is stronger than bond pair - bond pair electrons repulsion
hence caused the angle to squeeze to a smaller angle As
for water since there is a presence of lone pair - lone pair
electrons repulsion it results the bond pair - bond pair
electron repulsion to be much smaller hence caused the
angle to squeeze to a smaller angle Though valence bondtheory does not actually explain the hybridisation especially
when it relates to the repulsion occur involving lone pair
electron another theories shall be applied to study such
effect All these shall be discussed further in VSEPR theory
b) Effect of electronegativity and bonding angle between bond pair
- bond pair electrons ~
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- bond pair electrons
i As discussed earlier in the bonding of water (H2
O) the
molecular shape and angle is described in the diagram below
Sul hur which is also an element from Grou 16 formed h dro en
sulphide H2S when sulphur react with hydrogen However unlike
water the bonding angles is much smaller compare to water This is
due to the difference of electronegativity and also the bond length
between O and S in molecule Since O is more electronegative
compare to S the O-H bond is pulled closer toward O Furthermorethe bond length of O-H is shorter compare to S-H As a result the
bonding pair-bonding pair electrons repulse greater with each other in
H-O-H and caused the angle become greater
ii Another example is between NH3 and PH3 The orbital
f
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diagram for both NH3 and PH3 are described below Both
nitrogen N and phosphorous P are from the same groupwhich is Group 15
Similar to the above case P in phosphine is less
electronegative than N in ammonia and bond length of N-H
is shorter than P-H As a result H is pull closer to N and
repulsion between bonding of H-N-H is greater compare toH-P-H and caused the angle between H-N-H is greater
compare to H-P-H
324 Valence Shell Electron Pair Repulsion (VSEPR) Theory
1 M l l t i th th di i l t f t
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1 Molecular geometry is the three-dimensional arrangement of atoms
in a molecule A moleculersquos geometry affects its physical andchemical properties such as melting point boiling point density
and the types of reactions it undergoes The basic concept of
VSEPR theory is based on the three general rules below
a) As far as electron-pair repulsion is concerned double bonds andtriple bonds can be treated like single bonds This approximation is
good for qualitative purposes However you should realize that in
ldquo rdquo
because there are two or three bonds between two atoms theelectron density occupies more space
b) If a molecule has two or more resonance structures we can apply
the VSEPR model to any one of them Formal charges are usually
not shown
c) The order of repulsion strength of lone pair and bond pair are
lone-pair amp lone-pair electrons repulsion are the strongest
followed by lone-pair amp bond-pair electrons repulsion while
bond-pair amp bond-pair electrons repulsion is the weakest
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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bull Maximum overlap of bonding orbitals ~ The bond strength
depends on attraction of the nuclei for the shared electrons
so the greater the orbital overlap the stronger (morestable) the bond The extent of overlap depends on the
shapes and directions of the orbitals An s orbital is spherical
but p and d orbitals have more electron density in one
direction than in another Thus whenever possible a bond
involving p or d orbitals will be oriented in the direction that
maximizes overlap For example in hydrogen fluoride (HF)
bond the 1s orbital of H overlaps the half-filled 2p orbital of Falong the long axis of that orbital Any other direction would
result in less overlap and thus a weaker bond
H b idi ti f t i bit l T t f th
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bull Hybridisation of atomic orbitals ~ To account for the
bonding in simple diatomic molecules like HF we picture the
direct overlap of s and p orbitals of isolated atoms But howcan we account for the shapes of so many molecules and
polyatomic ions through the overlap of spherical s orbitals
dumbbell-shaped p orbitals and cloverleaf-shaped d orbitals
Linus Pauling proposed that the valence atomic orbitals in the
molecule are different from those in the isolated atoms The
spatial orientations of these new orbitals lead to more stable
bonds and are consistent with observed molecular shapesThe process of orbital mixing is called hybridisation and
the new atomic orbitals are called hybrid orbitals Two key
points about the number and type of hybrid orbitals are that
i The number of hybrid orbitals obtained equals the number ofatomic orbitals mixed
ii The type of hybrid orbitals obtained varies with the types of
atomic orbitals mixed
3 2 3 1 T f h b idi ti
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3231 Type of hybridisation
1 sp3 hybridisation ~ When four electron groups surround
the central atom the center atom involved must prepare fourorbitals with equal energies to overlap with the four
surrounding electron groups Valence Bond theory uses
hypothetical hybrid orbitals which are atomic orbitals
obtained when two or more non-equivalent orbitals of the
same atom combine in preparation for covalent bond
formation H bridisation is the term a lied to the mixin of
atomic orbitals in an atom (usually a central atom) togenerate a set of hybrid orbitals In the case of sp3 We can
generate four equivalent hybrid orbitals from the center
atom by mixing the s orbital and the three p orbitals
Explanation Energy level diagram Diagram of orbitals
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Explanation Energy level diagram Diagram of orbitals
a) Ground state
Carbon which act as thecenter atom has the valence
electron of 2s22p2
b) Excited state One of the electron from 2s
is promote to 2p orbitals -
equa energy eve
c) Hybridised state
One orbital of 2s and three
orbitals of 2p combined(hybrid) and rearrange
themselves to the shape
and orientation of a
tetrahedral shape
Molecular shape Tetrahedral
Angle between bond pair 10950
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2 sp2 hybridisation ~ When three electron groups surround
the central atom the center atom involved must prepare three
orbitals with equal energies to overlap with the three
surrounding electron groups In sp2 hybridisation three
equivalent (in terms of energy level) from the center atom by
mixing the one s orbital and the two p orbitals Using boron
trifluoride (BF3) as example sp2
hybridisation is explained
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3 sp hybridisation ~ When two electron groups surround
the central atom we observe a linear shape which
means that the bonding orbitals must have a linear
orientation VB theory explains this by proposing thatmixing two nonequivalent orbitals of a central atom one
s and one p gives rise to two equivalent sp hybrid
orbitals that lie 1800 apart
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Beryllium which act as the
center atom has the valence
electron of 2s2
b) Excited state
One of the electron from 2s ispromote to 2p orbitals - equal
energy level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of a
linear shape
Molecular shape Linear
Angle between bond pair 1800
4 The concept of hybridisation is also useful to explain molecules with
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p y p
doubletriple bonds By using the concept of the direct overlapping
orbitals and side-touch lapping orbitals the formation of multiple
bonds in ethene C2H4 and ethyne C2H2 are described
a) Ethene C2H4 - Hybridisation take place for both carbon atoms in
ethene molecule is sp2 hybridisation
bull When both hybridised C is bonded together one of the hybridisedorbital overlapped directly between each other while the other two
hybridised orbitals overlapped directly with hydrogen atoms
z
atoms form a side-touch bond between each other and form anotherbond as shown in the diagram below
bull From the diagram with C=C there are two types of bond A sigma-
bond (σ-bond) is covalent bonds formed by orbitals overlapping end-
to-end with the electron density concentrated between the nuclei ofthe bonding atoms while the second type is called a pi bond (π-
bond) which is defined as a covalent bond formed by sideways
overlapping orbitals with electron density concentrated above and
below the plane of the nuclei of the bonding atoms
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital of 2s and two
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of atrigonal planar shape
Note that on unhybridised pz
orbital an electron is presence
Molecular shape Trigonal planar Angle between bond pair 1200
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b) Ethyne C H has the Lewis structure of H CequivC H which the
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b) Ethyne C2H2 has the Lewis structure of HminusCequivCminusH which the
bonding can be explain using sp hybridisation Table below
described how the hybridisation take place on each carbonatom and how the formation of triple bond occur
c) From the diagram formation of -CequivC- is due to the formation
of one sigma-bond by direct overlapping of one of the twohybrid orbitals on each C while the other two pi-bond bonds
are formed as a result of side-lapping of the each two
y z
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital from 2s and 2p
orbitals combined (hybrid) and
rearrange themselves to the
shape and orientation of a
linear shapeNote that on unhybridised py amp
pz orbitals electrons are
presence to form two
π-bonds Molecular shape Linear Angle between bond pair 1800
sp
pz py
5 Other examples of applications in valence bond theory includes the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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formation of nitrogen molecule N2 and hydrogen cyanide HCN
molecule
a) Nitrogen gas is Earths most abundant gas as it cover 78 of the
content of our air Nitrogen molecule is an inert gas thanks to its
short covalent bond and also its strong triple bond Therefore
a lot of heats are required to break the chemical bond of nitrogenbefore it can be applied in industries Using valence bond theory
the bonding of nitrogen molecule is explained in the diagram
b) So when the two hybridised nitrogen atom interacting amongeach other they hence formed a linear shape and two pi-bonds
(πminusbonds) are formed as a result of side-lapping of unhybridised
py and pz orbital respectively
N N
ExplanationEnergy level
diagramDiagram of orbitals
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diagram
a) Ground state
Both nitrogen whichact as the center atom
has the valence
electron of 2s22p3
b) Excited state One of the electron
from 2s is promote to 2p
orbitals - e ual ener
level
c) Hybridised state
One orbital of 2s and
one orbitals of 2p
combined (hybrid) and
rearrange themselves
to the shape and
orientation of a linear
shapeMolecular shape Linear
Angle between bond pair 1800
b) Whereas for hydrogen cyanide HCN both carbon atom and
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) y g y
nitrogen atom undergoes sp hybridisation in order to form a
linear structureExplanation
Energy level diagram for
carbon atom C
Energy level diagram for
nitrogen atom N
a) Ground state
Valence electronC 2s22p2
N 2s22p3
b) Excited state
One of the electron from
2s is promote to 2p
orbitals - equal energy
level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined
(hybrid) and form sp
hybridisation
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N
σσσσ
π
σσσσ
π
6 However there are a few limitation on valence bond theory such
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as when explaining the effect angle of bond-pair and bond-pair
electrons when there isare presence of lone pair electron in thecenter atom and also the difference of electronegativity
a) Bonding in ammonia NH3 and water H2O ~ Both NH3 and
H2O undergoes an arrangement similar to sp3 hybridisation
similar to that of C in methane Table below shows thehybridisation occur for nitrogen in ammonia and oxygen in water
Energy level diagram Energy level diagram
for nitrogen atom N for oxygen atom O
a) Ground state
Valence electron
N 2s22p3
O 2s2
2p4
b) Excited state
One of the electron
from 2s is promote to
2p orbitals
c) Hybridised state
One orbital of 2s and
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One orbital of 2s and
three orbitals from 2p
combined (hybrid) andrearrange themselves to
the shape and
orientation similar to that
of tetrahedral shape
Similar to arrangement in
tetrahedral
Similar to arrangement in
tetrahedral
of ammonia NH3 andwater H2O
Shape and angleShape trigonal pyramidal
Angle 1070
Shape bent
Angle 10450
Number of bond pair amp
lone pair electrons
Bond pair electron 3
Lone pair electron 1
Bond pair electron 2
Lone pair electron 2
bull The angle of bond pair - bond pair electrons in ammonia and0 0
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water are 1070 and 10450 respectively which is lesser than in
methane molecule (10950
) This can be explained by the factof the presence of lone-pair electrons in both ammonia and
water Since the lone pair - lone pair electron repulsion is
stronger than lone pair - bond pair electron repulsion than
bond pair - bond pair electron repulsion it is expected that therepulsion between lone pair - bond pair electrons in ammonia
is stronger than bond pair - bond pair electrons repulsion
hence caused the angle to squeeze to a smaller angle As
for water since there is a presence of lone pair - lone pair
electrons repulsion it results the bond pair - bond pair
electron repulsion to be much smaller hence caused the
angle to squeeze to a smaller angle Though valence bondtheory does not actually explain the hybridisation especially
when it relates to the repulsion occur involving lone pair
electron another theories shall be applied to study such
effect All these shall be discussed further in VSEPR theory
b) Effect of electronegativity and bonding angle between bond pair
- bond pair electrons ~
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- bond pair electrons
i As discussed earlier in the bonding of water (H2
O) the
molecular shape and angle is described in the diagram below
Sul hur which is also an element from Grou 16 formed h dro en
sulphide H2S when sulphur react with hydrogen However unlike
water the bonding angles is much smaller compare to water This is
due to the difference of electronegativity and also the bond length
between O and S in molecule Since O is more electronegative
compare to S the O-H bond is pulled closer toward O Furthermorethe bond length of O-H is shorter compare to S-H As a result the
bonding pair-bonding pair electrons repulse greater with each other in
H-O-H and caused the angle become greater
ii Another example is between NH3 and PH3 The orbital
f
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diagram for both NH3 and PH3 are described below Both
nitrogen N and phosphorous P are from the same groupwhich is Group 15
Similar to the above case P in phosphine is less
electronegative than N in ammonia and bond length of N-H
is shorter than P-H As a result H is pull closer to N and
repulsion between bonding of H-N-H is greater compare toH-P-H and caused the angle between H-N-H is greater
compare to H-P-H
324 Valence Shell Electron Pair Repulsion (VSEPR) Theory
1 M l l t i th th di i l t f t
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1 Molecular geometry is the three-dimensional arrangement of atoms
in a molecule A moleculersquos geometry affects its physical andchemical properties such as melting point boiling point density
and the types of reactions it undergoes The basic concept of
VSEPR theory is based on the three general rules below
a) As far as electron-pair repulsion is concerned double bonds andtriple bonds can be treated like single bonds This approximation is
good for qualitative purposes However you should realize that in
ldquo rdquo
because there are two or three bonds between two atoms theelectron density occupies more space
b) If a molecule has two or more resonance structures we can apply
the VSEPR model to any one of them Formal charges are usually
not shown
c) The order of repulsion strength of lone pair and bond pair are
lone-pair amp lone-pair electrons repulsion are the strongest
followed by lone-pair amp bond-pair electrons repulsion while
bond-pair amp bond-pair electrons repulsion is the weakest
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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bull Hybridisation of atomic orbitals ~ To account for the
bonding in simple diatomic molecules like HF we picture the
direct overlap of s and p orbitals of isolated atoms But howcan we account for the shapes of so many molecules and
polyatomic ions through the overlap of spherical s orbitals
dumbbell-shaped p orbitals and cloverleaf-shaped d orbitals
Linus Pauling proposed that the valence atomic orbitals in the
molecule are different from those in the isolated atoms The
spatial orientations of these new orbitals lead to more stable
bonds and are consistent with observed molecular shapesThe process of orbital mixing is called hybridisation and
the new atomic orbitals are called hybrid orbitals Two key
points about the number and type of hybrid orbitals are that
i The number of hybrid orbitals obtained equals the number ofatomic orbitals mixed
ii The type of hybrid orbitals obtained varies with the types of
atomic orbitals mixed
3 2 3 1 T f h b idi ti
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3231 Type of hybridisation
1 sp3 hybridisation ~ When four electron groups surround
the central atom the center atom involved must prepare fourorbitals with equal energies to overlap with the four
surrounding electron groups Valence Bond theory uses
hypothetical hybrid orbitals which are atomic orbitals
obtained when two or more non-equivalent orbitals of the
same atom combine in preparation for covalent bond
formation H bridisation is the term a lied to the mixin of
atomic orbitals in an atom (usually a central atom) togenerate a set of hybrid orbitals In the case of sp3 We can
generate four equivalent hybrid orbitals from the center
atom by mixing the s orbital and the three p orbitals
Explanation Energy level diagram Diagram of orbitals
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Explanation Energy level diagram Diagram of orbitals
a) Ground state
Carbon which act as thecenter atom has the valence
electron of 2s22p2
b) Excited state One of the electron from 2s
is promote to 2p orbitals -
equa energy eve
c) Hybridised state
One orbital of 2s and three
orbitals of 2p combined(hybrid) and rearrange
themselves to the shape
and orientation of a
tetrahedral shape
Molecular shape Tetrahedral
Angle between bond pair 10950
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2 sp2 hybridisation ~ When three electron groups surround
the central atom the center atom involved must prepare three
orbitals with equal energies to overlap with the three
surrounding electron groups In sp2 hybridisation three
equivalent (in terms of energy level) from the center atom by
mixing the one s orbital and the two p orbitals Using boron
trifluoride (BF3) as example sp2
hybridisation is explained
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3 sp hybridisation ~ When two electron groups surround
the central atom we observe a linear shape which
means that the bonding orbitals must have a linear
orientation VB theory explains this by proposing thatmixing two nonequivalent orbitals of a central atom one
s and one p gives rise to two equivalent sp hybrid
orbitals that lie 1800 apart
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Beryllium which act as the
center atom has the valence
electron of 2s2
b) Excited state
One of the electron from 2s ispromote to 2p orbitals - equal
energy level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of a
linear shape
Molecular shape Linear
Angle between bond pair 1800
4 The concept of hybridisation is also useful to explain molecules with
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p y p
doubletriple bonds By using the concept of the direct overlapping
orbitals and side-touch lapping orbitals the formation of multiple
bonds in ethene C2H4 and ethyne C2H2 are described
a) Ethene C2H4 - Hybridisation take place for both carbon atoms in
ethene molecule is sp2 hybridisation
bull When both hybridised C is bonded together one of the hybridisedorbital overlapped directly between each other while the other two
hybridised orbitals overlapped directly with hydrogen atoms
z
atoms form a side-touch bond between each other and form anotherbond as shown in the diagram below
bull From the diagram with C=C there are two types of bond A sigma-
bond (σ-bond) is covalent bonds formed by orbitals overlapping end-
to-end with the electron density concentrated between the nuclei ofthe bonding atoms while the second type is called a pi bond (π-
bond) which is defined as a covalent bond formed by sideways
overlapping orbitals with electron density concentrated above and
below the plane of the nuclei of the bonding atoms
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital of 2s and two
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of atrigonal planar shape
Note that on unhybridised pz
orbital an electron is presence
Molecular shape Trigonal planar Angle between bond pair 1200
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b) Ethyne C H has the Lewis structure of H CequivC H which the
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b) Ethyne C2H2 has the Lewis structure of HminusCequivCminusH which the
bonding can be explain using sp hybridisation Table below
described how the hybridisation take place on each carbonatom and how the formation of triple bond occur
c) From the diagram formation of -CequivC- is due to the formation
of one sigma-bond by direct overlapping of one of the twohybrid orbitals on each C while the other two pi-bond bonds
are formed as a result of side-lapping of the each two
y z
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital from 2s and 2p
orbitals combined (hybrid) and
rearrange themselves to the
shape and orientation of a
linear shapeNote that on unhybridised py amp
pz orbitals electrons are
presence to form two
π-bonds Molecular shape Linear Angle between bond pair 1800
sp
pz py
5 Other examples of applications in valence bond theory includes the
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formation of nitrogen molecule N2 and hydrogen cyanide HCN
molecule
a) Nitrogen gas is Earths most abundant gas as it cover 78 of the
content of our air Nitrogen molecule is an inert gas thanks to its
short covalent bond and also its strong triple bond Therefore
a lot of heats are required to break the chemical bond of nitrogenbefore it can be applied in industries Using valence bond theory
the bonding of nitrogen molecule is explained in the diagram
b) So when the two hybridised nitrogen atom interacting amongeach other they hence formed a linear shape and two pi-bonds
(πminusbonds) are formed as a result of side-lapping of unhybridised
py and pz orbital respectively
N N
ExplanationEnergy level
diagramDiagram of orbitals
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diagram
a) Ground state
Both nitrogen whichact as the center atom
has the valence
electron of 2s22p3
b) Excited state One of the electron
from 2s is promote to 2p
orbitals - e ual ener
level
c) Hybridised state
One orbital of 2s and
one orbitals of 2p
combined (hybrid) and
rearrange themselves
to the shape and
orientation of a linear
shapeMolecular shape Linear
Angle between bond pair 1800
b) Whereas for hydrogen cyanide HCN both carbon atom and
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) y g y
nitrogen atom undergoes sp hybridisation in order to form a
linear structureExplanation
Energy level diagram for
carbon atom C
Energy level diagram for
nitrogen atom N
a) Ground state
Valence electronC 2s22p2
N 2s22p3
b) Excited state
One of the electron from
2s is promote to 2p
orbitals - equal energy
level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined
(hybrid) and form sp
hybridisation
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N
σσσσ
π
σσσσ
π
6 However there are a few limitation on valence bond theory such
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as when explaining the effect angle of bond-pair and bond-pair
electrons when there isare presence of lone pair electron in thecenter atom and also the difference of electronegativity
a) Bonding in ammonia NH3 and water H2O ~ Both NH3 and
H2O undergoes an arrangement similar to sp3 hybridisation
similar to that of C in methane Table below shows thehybridisation occur for nitrogen in ammonia and oxygen in water
Energy level diagram Energy level diagram
for nitrogen atom N for oxygen atom O
a) Ground state
Valence electron
N 2s22p3
O 2s2
2p4
b) Excited state
One of the electron
from 2s is promote to
2p orbitals
c) Hybridised state
One orbital of 2s and
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One orbital of 2s and
three orbitals from 2p
combined (hybrid) andrearrange themselves to
the shape and
orientation similar to that
of tetrahedral shape
Similar to arrangement in
tetrahedral
Similar to arrangement in
tetrahedral
of ammonia NH3 andwater H2O
Shape and angleShape trigonal pyramidal
Angle 1070
Shape bent
Angle 10450
Number of bond pair amp
lone pair electrons
Bond pair electron 3
Lone pair electron 1
Bond pair electron 2
Lone pair electron 2
bull The angle of bond pair - bond pair electrons in ammonia and0 0
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water are 1070 and 10450 respectively which is lesser than in
methane molecule (10950
) This can be explained by the factof the presence of lone-pair electrons in both ammonia and
water Since the lone pair - lone pair electron repulsion is
stronger than lone pair - bond pair electron repulsion than
bond pair - bond pair electron repulsion it is expected that therepulsion between lone pair - bond pair electrons in ammonia
is stronger than bond pair - bond pair electrons repulsion
hence caused the angle to squeeze to a smaller angle As
for water since there is a presence of lone pair - lone pair
electrons repulsion it results the bond pair - bond pair
electron repulsion to be much smaller hence caused the
angle to squeeze to a smaller angle Though valence bondtheory does not actually explain the hybridisation especially
when it relates to the repulsion occur involving lone pair
electron another theories shall be applied to study such
effect All these shall be discussed further in VSEPR theory
b) Effect of electronegativity and bonding angle between bond pair
- bond pair electrons ~
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- bond pair electrons
i As discussed earlier in the bonding of water (H2
O) the
molecular shape and angle is described in the diagram below
Sul hur which is also an element from Grou 16 formed h dro en
sulphide H2S when sulphur react with hydrogen However unlike
water the bonding angles is much smaller compare to water This is
due to the difference of electronegativity and also the bond length
between O and S in molecule Since O is more electronegative
compare to S the O-H bond is pulled closer toward O Furthermorethe bond length of O-H is shorter compare to S-H As a result the
bonding pair-bonding pair electrons repulse greater with each other in
H-O-H and caused the angle become greater
ii Another example is between NH3 and PH3 The orbital
f
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diagram for both NH3 and PH3 are described below Both
nitrogen N and phosphorous P are from the same groupwhich is Group 15
Similar to the above case P in phosphine is less
electronegative than N in ammonia and bond length of N-H
is shorter than P-H As a result H is pull closer to N and
repulsion between bonding of H-N-H is greater compare toH-P-H and caused the angle between H-N-H is greater
compare to H-P-H
324 Valence Shell Electron Pair Repulsion (VSEPR) Theory
1 M l l t i th th di i l t f t
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1 Molecular geometry is the three-dimensional arrangement of atoms
in a molecule A moleculersquos geometry affects its physical andchemical properties such as melting point boiling point density
and the types of reactions it undergoes The basic concept of
VSEPR theory is based on the three general rules below
a) As far as electron-pair repulsion is concerned double bonds andtriple bonds can be treated like single bonds This approximation is
good for qualitative purposes However you should realize that in
ldquo rdquo
because there are two or three bonds between two atoms theelectron density occupies more space
b) If a molecule has two or more resonance structures we can apply
the VSEPR model to any one of them Formal charges are usually
not shown
c) The order of repulsion strength of lone pair and bond pair are
lone-pair amp lone-pair electrons repulsion are the strongest
followed by lone-pair amp bond-pair electrons repulsion while
bond-pair amp bond-pair electrons repulsion is the weakest
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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3231 Type of hybridisation
1 sp3 hybridisation ~ When four electron groups surround
the central atom the center atom involved must prepare fourorbitals with equal energies to overlap with the four
surrounding electron groups Valence Bond theory uses
hypothetical hybrid orbitals which are atomic orbitals
obtained when two or more non-equivalent orbitals of the
same atom combine in preparation for covalent bond
formation H bridisation is the term a lied to the mixin of
atomic orbitals in an atom (usually a central atom) togenerate a set of hybrid orbitals In the case of sp3 We can
generate four equivalent hybrid orbitals from the center
atom by mixing the s orbital and the three p orbitals
Explanation Energy level diagram Diagram of orbitals
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Explanation Energy level diagram Diagram of orbitals
a) Ground state
Carbon which act as thecenter atom has the valence
electron of 2s22p2
b) Excited state One of the electron from 2s
is promote to 2p orbitals -
equa energy eve
c) Hybridised state
One orbital of 2s and three
orbitals of 2p combined(hybrid) and rearrange
themselves to the shape
and orientation of a
tetrahedral shape
Molecular shape Tetrahedral
Angle between bond pair 10950
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2 sp2 hybridisation ~ When three electron groups surround
the central atom the center atom involved must prepare three
orbitals with equal energies to overlap with the three
surrounding electron groups In sp2 hybridisation three
equivalent (in terms of energy level) from the center atom by
mixing the one s orbital and the two p orbitals Using boron
trifluoride (BF3) as example sp2
hybridisation is explained
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3 sp hybridisation ~ When two electron groups surround
the central atom we observe a linear shape which
means that the bonding orbitals must have a linear
orientation VB theory explains this by proposing thatmixing two nonequivalent orbitals of a central atom one
s and one p gives rise to two equivalent sp hybrid
orbitals that lie 1800 apart
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Beryllium which act as the
center atom has the valence
electron of 2s2
b) Excited state
One of the electron from 2s ispromote to 2p orbitals - equal
energy level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of a
linear shape
Molecular shape Linear
Angle between bond pair 1800
4 The concept of hybridisation is also useful to explain molecules with
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p y p
doubletriple bonds By using the concept of the direct overlapping
orbitals and side-touch lapping orbitals the formation of multiple
bonds in ethene C2H4 and ethyne C2H2 are described
a) Ethene C2H4 - Hybridisation take place for both carbon atoms in
ethene molecule is sp2 hybridisation
bull When both hybridised C is bonded together one of the hybridisedorbital overlapped directly between each other while the other two
hybridised orbitals overlapped directly with hydrogen atoms
z
atoms form a side-touch bond between each other and form anotherbond as shown in the diagram below
bull From the diagram with C=C there are two types of bond A sigma-
bond (σ-bond) is covalent bonds formed by orbitals overlapping end-
to-end with the electron density concentrated between the nuclei ofthe bonding atoms while the second type is called a pi bond (π-
bond) which is defined as a covalent bond formed by sideways
overlapping orbitals with electron density concentrated above and
below the plane of the nuclei of the bonding atoms
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital of 2s and two
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of atrigonal planar shape
Note that on unhybridised pz
orbital an electron is presence
Molecular shape Trigonal planar Angle between bond pair 1200
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b) Ethyne C H has the Lewis structure of H CequivC H which the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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b) Ethyne C2H2 has the Lewis structure of HminusCequivCminusH which the
bonding can be explain using sp hybridisation Table below
described how the hybridisation take place on each carbonatom and how the formation of triple bond occur
c) From the diagram formation of -CequivC- is due to the formation
of one sigma-bond by direct overlapping of one of the twohybrid orbitals on each C while the other two pi-bond bonds
are formed as a result of side-lapping of the each two
y z
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital from 2s and 2p
orbitals combined (hybrid) and
rearrange themselves to the
shape and orientation of a
linear shapeNote that on unhybridised py amp
pz orbitals electrons are
presence to form two
π-bonds Molecular shape Linear Angle between bond pair 1800
sp
pz py
5 Other examples of applications in valence bond theory includes the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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formation of nitrogen molecule N2 and hydrogen cyanide HCN
molecule
a) Nitrogen gas is Earths most abundant gas as it cover 78 of the
content of our air Nitrogen molecule is an inert gas thanks to its
short covalent bond and also its strong triple bond Therefore
a lot of heats are required to break the chemical bond of nitrogenbefore it can be applied in industries Using valence bond theory
the bonding of nitrogen molecule is explained in the diagram
b) So when the two hybridised nitrogen atom interacting amongeach other they hence formed a linear shape and two pi-bonds
(πminusbonds) are formed as a result of side-lapping of unhybridised
py and pz orbital respectively
N N
ExplanationEnergy level
diagramDiagram of orbitals
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diagram
a) Ground state
Both nitrogen whichact as the center atom
has the valence
electron of 2s22p3
b) Excited state One of the electron
from 2s is promote to 2p
orbitals - e ual ener
level
c) Hybridised state
One orbital of 2s and
one orbitals of 2p
combined (hybrid) and
rearrange themselves
to the shape and
orientation of a linear
shapeMolecular shape Linear
Angle between bond pair 1800
b) Whereas for hydrogen cyanide HCN both carbon atom and
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) y g y
nitrogen atom undergoes sp hybridisation in order to form a
linear structureExplanation
Energy level diagram for
carbon atom C
Energy level diagram for
nitrogen atom N
a) Ground state
Valence electronC 2s22p2
N 2s22p3
b) Excited state
One of the electron from
2s is promote to 2p
orbitals - equal energy
level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined
(hybrid) and form sp
hybridisation
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N
σσσσ
π
σσσσ
π
6 However there are a few limitation on valence bond theory such
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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as when explaining the effect angle of bond-pair and bond-pair
electrons when there isare presence of lone pair electron in thecenter atom and also the difference of electronegativity
a) Bonding in ammonia NH3 and water H2O ~ Both NH3 and
H2O undergoes an arrangement similar to sp3 hybridisation
similar to that of C in methane Table below shows thehybridisation occur for nitrogen in ammonia and oxygen in water
Energy level diagram Energy level diagram
for nitrogen atom N for oxygen atom O
a) Ground state
Valence electron
N 2s22p3
O 2s2
2p4
b) Excited state
One of the electron
from 2s is promote to
2p orbitals
c) Hybridised state
One orbital of 2s and
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One orbital of 2s and
three orbitals from 2p
combined (hybrid) andrearrange themselves to
the shape and
orientation similar to that
of tetrahedral shape
Similar to arrangement in
tetrahedral
Similar to arrangement in
tetrahedral
of ammonia NH3 andwater H2O
Shape and angleShape trigonal pyramidal
Angle 1070
Shape bent
Angle 10450
Number of bond pair amp
lone pair electrons
Bond pair electron 3
Lone pair electron 1
Bond pair electron 2
Lone pair electron 2
bull The angle of bond pair - bond pair electrons in ammonia and0 0
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water are 1070 and 10450 respectively which is lesser than in
methane molecule (10950
) This can be explained by the factof the presence of lone-pair electrons in both ammonia and
water Since the lone pair - lone pair electron repulsion is
stronger than lone pair - bond pair electron repulsion than
bond pair - bond pair electron repulsion it is expected that therepulsion between lone pair - bond pair electrons in ammonia
is stronger than bond pair - bond pair electrons repulsion
hence caused the angle to squeeze to a smaller angle As
for water since there is a presence of lone pair - lone pair
electrons repulsion it results the bond pair - bond pair
electron repulsion to be much smaller hence caused the
angle to squeeze to a smaller angle Though valence bondtheory does not actually explain the hybridisation especially
when it relates to the repulsion occur involving lone pair
electron another theories shall be applied to study such
effect All these shall be discussed further in VSEPR theory
b) Effect of electronegativity and bonding angle between bond pair
- bond pair electrons ~
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- bond pair electrons
i As discussed earlier in the bonding of water (H2
O) the
molecular shape and angle is described in the diagram below
Sul hur which is also an element from Grou 16 formed h dro en
sulphide H2S when sulphur react with hydrogen However unlike
water the bonding angles is much smaller compare to water This is
due to the difference of electronegativity and also the bond length
between O and S in molecule Since O is more electronegative
compare to S the O-H bond is pulled closer toward O Furthermorethe bond length of O-H is shorter compare to S-H As a result the
bonding pair-bonding pair electrons repulse greater with each other in
H-O-H and caused the angle become greater
ii Another example is between NH3 and PH3 The orbital
f
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diagram for both NH3 and PH3 are described below Both
nitrogen N and phosphorous P are from the same groupwhich is Group 15
Similar to the above case P in phosphine is less
electronegative than N in ammonia and bond length of N-H
is shorter than P-H As a result H is pull closer to N and
repulsion between bonding of H-N-H is greater compare toH-P-H and caused the angle between H-N-H is greater
compare to H-P-H
324 Valence Shell Electron Pair Repulsion (VSEPR) Theory
1 M l l t i th th di i l t f t
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1 Molecular geometry is the three-dimensional arrangement of atoms
in a molecule A moleculersquos geometry affects its physical andchemical properties such as melting point boiling point density
and the types of reactions it undergoes The basic concept of
VSEPR theory is based on the three general rules below
a) As far as electron-pair repulsion is concerned double bonds andtriple bonds can be treated like single bonds This approximation is
good for qualitative purposes However you should realize that in
ldquo rdquo
because there are two or three bonds between two atoms theelectron density occupies more space
b) If a molecule has two or more resonance structures we can apply
the VSEPR model to any one of them Formal charges are usually
not shown
c) The order of repulsion strength of lone pair and bond pair are
lone-pair amp lone-pair electrons repulsion are the strongest
followed by lone-pair amp bond-pair electrons repulsion while
bond-pair amp bond-pair electrons repulsion is the weakest
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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Explanation Energy level diagram Diagram of orbitals
a) Ground state
Carbon which act as thecenter atom has the valence
electron of 2s22p2
b) Excited state One of the electron from 2s
is promote to 2p orbitals -
equa energy eve
c) Hybridised state
One orbital of 2s and three
orbitals of 2p combined(hybrid) and rearrange
themselves to the shape
and orientation of a
tetrahedral shape
Molecular shape Tetrahedral
Angle between bond pair 10950
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2 sp2 hybridisation ~ When three electron groups surround
the central atom the center atom involved must prepare three
orbitals with equal energies to overlap with the three
surrounding electron groups In sp2 hybridisation three
equivalent (in terms of energy level) from the center atom by
mixing the one s orbital and the two p orbitals Using boron
trifluoride (BF3) as example sp2
hybridisation is explained
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3 sp hybridisation ~ When two electron groups surround
the central atom we observe a linear shape which
means that the bonding orbitals must have a linear
orientation VB theory explains this by proposing thatmixing two nonequivalent orbitals of a central atom one
s and one p gives rise to two equivalent sp hybrid
orbitals that lie 1800 apart
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Beryllium which act as the
center atom has the valence
electron of 2s2
b) Excited state
One of the electron from 2s ispromote to 2p orbitals - equal
energy level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of a
linear shape
Molecular shape Linear
Angle between bond pair 1800
4 The concept of hybridisation is also useful to explain molecules with
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p y p
doubletriple bonds By using the concept of the direct overlapping
orbitals and side-touch lapping orbitals the formation of multiple
bonds in ethene C2H4 and ethyne C2H2 are described
a) Ethene C2H4 - Hybridisation take place for both carbon atoms in
ethene molecule is sp2 hybridisation
bull When both hybridised C is bonded together one of the hybridisedorbital overlapped directly between each other while the other two
hybridised orbitals overlapped directly with hydrogen atoms
z
atoms form a side-touch bond between each other and form anotherbond as shown in the diagram below
bull From the diagram with C=C there are two types of bond A sigma-
bond (σ-bond) is covalent bonds formed by orbitals overlapping end-
to-end with the electron density concentrated between the nuclei ofthe bonding atoms while the second type is called a pi bond (π-
bond) which is defined as a covalent bond formed by sideways
overlapping orbitals with electron density concentrated above and
below the plane of the nuclei of the bonding atoms
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital of 2s and two
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of atrigonal planar shape
Note that on unhybridised pz
orbital an electron is presence
Molecular shape Trigonal planar Angle between bond pair 1200
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b) Ethyne C H has the Lewis structure of H CequivC H which the
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b) Ethyne C2H2 has the Lewis structure of HminusCequivCminusH which the
bonding can be explain using sp hybridisation Table below
described how the hybridisation take place on each carbonatom and how the formation of triple bond occur
c) From the diagram formation of -CequivC- is due to the formation
of one sigma-bond by direct overlapping of one of the twohybrid orbitals on each C while the other two pi-bond bonds
are formed as a result of side-lapping of the each two
y z
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital from 2s and 2p
orbitals combined (hybrid) and
rearrange themselves to the
shape and orientation of a
linear shapeNote that on unhybridised py amp
pz orbitals electrons are
presence to form two
π-bonds Molecular shape Linear Angle between bond pair 1800
sp
pz py
5 Other examples of applications in valence bond theory includes the
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formation of nitrogen molecule N2 and hydrogen cyanide HCN
molecule
a) Nitrogen gas is Earths most abundant gas as it cover 78 of the
content of our air Nitrogen molecule is an inert gas thanks to its
short covalent bond and also its strong triple bond Therefore
a lot of heats are required to break the chemical bond of nitrogenbefore it can be applied in industries Using valence bond theory
the bonding of nitrogen molecule is explained in the diagram
b) So when the two hybridised nitrogen atom interacting amongeach other they hence formed a linear shape and two pi-bonds
(πminusbonds) are formed as a result of side-lapping of unhybridised
py and pz orbital respectively
N N
ExplanationEnergy level
diagramDiagram of orbitals
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diagram
a) Ground state
Both nitrogen whichact as the center atom
has the valence
electron of 2s22p3
b) Excited state One of the electron
from 2s is promote to 2p
orbitals - e ual ener
level
c) Hybridised state
One orbital of 2s and
one orbitals of 2p
combined (hybrid) and
rearrange themselves
to the shape and
orientation of a linear
shapeMolecular shape Linear
Angle between bond pair 1800
b) Whereas for hydrogen cyanide HCN both carbon atom and
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) y g y
nitrogen atom undergoes sp hybridisation in order to form a
linear structureExplanation
Energy level diagram for
carbon atom C
Energy level diagram for
nitrogen atom N
a) Ground state
Valence electronC 2s22p2
N 2s22p3
b) Excited state
One of the electron from
2s is promote to 2p
orbitals - equal energy
level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined
(hybrid) and form sp
hybridisation
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N
σσσσ
π
σσσσ
π
6 However there are a few limitation on valence bond theory such
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as when explaining the effect angle of bond-pair and bond-pair
electrons when there isare presence of lone pair electron in thecenter atom and also the difference of electronegativity
a) Bonding in ammonia NH3 and water H2O ~ Both NH3 and
H2O undergoes an arrangement similar to sp3 hybridisation
similar to that of C in methane Table below shows thehybridisation occur for nitrogen in ammonia and oxygen in water
Energy level diagram Energy level diagram
for nitrogen atom N for oxygen atom O
a) Ground state
Valence electron
N 2s22p3
O 2s2
2p4
b) Excited state
One of the electron
from 2s is promote to
2p orbitals
c) Hybridised state
One orbital of 2s and
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One orbital of 2s and
three orbitals from 2p
combined (hybrid) andrearrange themselves to
the shape and
orientation similar to that
of tetrahedral shape
Similar to arrangement in
tetrahedral
Similar to arrangement in
tetrahedral
of ammonia NH3 andwater H2O
Shape and angleShape trigonal pyramidal
Angle 1070
Shape bent
Angle 10450
Number of bond pair amp
lone pair electrons
Bond pair electron 3
Lone pair electron 1
Bond pair electron 2
Lone pair electron 2
bull The angle of bond pair - bond pair electrons in ammonia and0 0
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water are 1070 and 10450 respectively which is lesser than in
methane molecule (10950
) This can be explained by the factof the presence of lone-pair electrons in both ammonia and
water Since the lone pair - lone pair electron repulsion is
stronger than lone pair - bond pair electron repulsion than
bond pair - bond pair electron repulsion it is expected that therepulsion between lone pair - bond pair electrons in ammonia
is stronger than bond pair - bond pair electrons repulsion
hence caused the angle to squeeze to a smaller angle As
for water since there is a presence of lone pair - lone pair
electrons repulsion it results the bond pair - bond pair
electron repulsion to be much smaller hence caused the
angle to squeeze to a smaller angle Though valence bondtheory does not actually explain the hybridisation especially
when it relates to the repulsion occur involving lone pair
electron another theories shall be applied to study such
effect All these shall be discussed further in VSEPR theory
b) Effect of electronegativity and bonding angle between bond pair
- bond pair electrons ~
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- bond pair electrons
i As discussed earlier in the bonding of water (H2
O) the
molecular shape and angle is described in the diagram below
Sul hur which is also an element from Grou 16 formed h dro en
sulphide H2S when sulphur react with hydrogen However unlike
water the bonding angles is much smaller compare to water This is
due to the difference of electronegativity and also the bond length
between O and S in molecule Since O is more electronegative
compare to S the O-H bond is pulled closer toward O Furthermorethe bond length of O-H is shorter compare to S-H As a result the
bonding pair-bonding pair electrons repulse greater with each other in
H-O-H and caused the angle become greater
ii Another example is between NH3 and PH3 The orbital
f
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diagram for both NH3 and PH3 are described below Both
nitrogen N and phosphorous P are from the same groupwhich is Group 15
Similar to the above case P in phosphine is less
electronegative than N in ammonia and bond length of N-H
is shorter than P-H As a result H is pull closer to N and
repulsion between bonding of H-N-H is greater compare toH-P-H and caused the angle between H-N-H is greater
compare to H-P-H
324 Valence Shell Electron Pair Repulsion (VSEPR) Theory
1 M l l t i th th di i l t f t
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1 Molecular geometry is the three-dimensional arrangement of atoms
in a molecule A moleculersquos geometry affects its physical andchemical properties such as melting point boiling point density
and the types of reactions it undergoes The basic concept of
VSEPR theory is based on the three general rules below
a) As far as electron-pair repulsion is concerned double bonds andtriple bonds can be treated like single bonds This approximation is
good for qualitative purposes However you should realize that in
ldquo rdquo
because there are two or three bonds between two atoms theelectron density occupies more space
b) If a molecule has two or more resonance structures we can apply
the VSEPR model to any one of them Formal charges are usually
not shown
c) The order of repulsion strength of lone pair and bond pair are
lone-pair amp lone-pair electrons repulsion are the strongest
followed by lone-pair amp bond-pair electrons repulsion while
bond-pair amp bond-pair electrons repulsion is the weakest
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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2 sp2 hybridisation ~ When three electron groups surround
the central atom the center atom involved must prepare three
orbitals with equal energies to overlap with the three
surrounding electron groups In sp2 hybridisation three
equivalent (in terms of energy level) from the center atom by
mixing the one s orbital and the two p orbitals Using boron
trifluoride (BF3) as example sp2
hybridisation is explained
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8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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3 sp hybridisation ~ When two electron groups surround
the central atom we observe a linear shape which
means that the bonding orbitals must have a linear
orientation VB theory explains this by proposing thatmixing two nonequivalent orbitals of a central atom one
s and one p gives rise to two equivalent sp hybrid
orbitals that lie 1800 apart
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Beryllium which act as the
center atom has the valence
electron of 2s2
b) Excited state
One of the electron from 2s ispromote to 2p orbitals - equal
energy level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of a
linear shape
Molecular shape Linear
Angle between bond pair 1800
4 The concept of hybridisation is also useful to explain molecules with
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p y p
doubletriple bonds By using the concept of the direct overlapping
orbitals and side-touch lapping orbitals the formation of multiple
bonds in ethene C2H4 and ethyne C2H2 are described
a) Ethene C2H4 - Hybridisation take place for both carbon atoms in
ethene molecule is sp2 hybridisation
bull When both hybridised C is bonded together one of the hybridisedorbital overlapped directly between each other while the other two
hybridised orbitals overlapped directly with hydrogen atoms
z
atoms form a side-touch bond between each other and form anotherbond as shown in the diagram below
bull From the diagram with C=C there are two types of bond A sigma-
bond (σ-bond) is covalent bonds formed by orbitals overlapping end-
to-end with the electron density concentrated between the nuclei ofthe bonding atoms while the second type is called a pi bond (π-
bond) which is defined as a covalent bond formed by sideways
overlapping orbitals with electron density concentrated above and
below the plane of the nuclei of the bonding atoms
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital of 2s and two
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of atrigonal planar shape
Note that on unhybridised pz
orbital an electron is presence
Molecular shape Trigonal planar Angle between bond pair 1200
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b) Ethyne C H has the Lewis structure of H CequivC H which the
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b) Ethyne C2H2 has the Lewis structure of HminusCequivCminusH which the
bonding can be explain using sp hybridisation Table below
described how the hybridisation take place on each carbonatom and how the formation of triple bond occur
c) From the diagram formation of -CequivC- is due to the formation
of one sigma-bond by direct overlapping of one of the twohybrid orbitals on each C while the other two pi-bond bonds
are formed as a result of side-lapping of the each two
y z
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital from 2s and 2p
orbitals combined (hybrid) and
rearrange themselves to the
shape and orientation of a
linear shapeNote that on unhybridised py amp
pz orbitals electrons are
presence to form two
π-bonds Molecular shape Linear Angle between bond pair 1800
sp
pz py
5 Other examples of applications in valence bond theory includes the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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formation of nitrogen molecule N2 and hydrogen cyanide HCN
molecule
a) Nitrogen gas is Earths most abundant gas as it cover 78 of the
content of our air Nitrogen molecule is an inert gas thanks to its
short covalent bond and also its strong triple bond Therefore
a lot of heats are required to break the chemical bond of nitrogenbefore it can be applied in industries Using valence bond theory
the bonding of nitrogen molecule is explained in the diagram
b) So when the two hybridised nitrogen atom interacting amongeach other they hence formed a linear shape and two pi-bonds
(πminusbonds) are formed as a result of side-lapping of unhybridised
py and pz orbital respectively
N N
ExplanationEnergy level
diagramDiagram of orbitals
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diagram
a) Ground state
Both nitrogen whichact as the center atom
has the valence
electron of 2s22p3
b) Excited state One of the electron
from 2s is promote to 2p
orbitals - e ual ener
level
c) Hybridised state
One orbital of 2s and
one orbitals of 2p
combined (hybrid) and
rearrange themselves
to the shape and
orientation of a linear
shapeMolecular shape Linear
Angle between bond pair 1800
b) Whereas for hydrogen cyanide HCN both carbon atom and
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) y g y
nitrogen atom undergoes sp hybridisation in order to form a
linear structureExplanation
Energy level diagram for
carbon atom C
Energy level diagram for
nitrogen atom N
a) Ground state
Valence electronC 2s22p2
N 2s22p3
b) Excited state
One of the electron from
2s is promote to 2p
orbitals - equal energy
level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined
(hybrid) and form sp
hybridisation
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N
σσσσ
π
σσσσ
π
6 However there are a few limitation on valence bond theory such
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as when explaining the effect angle of bond-pair and bond-pair
electrons when there isare presence of lone pair electron in thecenter atom and also the difference of electronegativity
a) Bonding in ammonia NH3 and water H2O ~ Both NH3 and
H2O undergoes an arrangement similar to sp3 hybridisation
similar to that of C in methane Table below shows thehybridisation occur for nitrogen in ammonia and oxygen in water
Energy level diagram Energy level diagram
for nitrogen atom N for oxygen atom O
a) Ground state
Valence electron
N 2s22p3
O 2s2
2p4
b) Excited state
One of the electron
from 2s is promote to
2p orbitals
c) Hybridised state
One orbital of 2s and
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One orbital of 2s and
three orbitals from 2p
combined (hybrid) andrearrange themselves to
the shape and
orientation similar to that
of tetrahedral shape
Similar to arrangement in
tetrahedral
Similar to arrangement in
tetrahedral
of ammonia NH3 andwater H2O
Shape and angleShape trigonal pyramidal
Angle 1070
Shape bent
Angle 10450
Number of bond pair amp
lone pair electrons
Bond pair electron 3
Lone pair electron 1
Bond pair electron 2
Lone pair electron 2
bull The angle of bond pair - bond pair electrons in ammonia and0 0
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water are 1070 and 10450 respectively which is lesser than in
methane molecule (10950
) This can be explained by the factof the presence of lone-pair electrons in both ammonia and
water Since the lone pair - lone pair electron repulsion is
stronger than lone pair - bond pair electron repulsion than
bond pair - bond pair electron repulsion it is expected that therepulsion between lone pair - bond pair electrons in ammonia
is stronger than bond pair - bond pair electrons repulsion
hence caused the angle to squeeze to a smaller angle As
for water since there is a presence of lone pair - lone pair
electrons repulsion it results the bond pair - bond pair
electron repulsion to be much smaller hence caused the
angle to squeeze to a smaller angle Though valence bondtheory does not actually explain the hybridisation especially
when it relates to the repulsion occur involving lone pair
electron another theories shall be applied to study such
effect All these shall be discussed further in VSEPR theory
b) Effect of electronegativity and bonding angle between bond pair
- bond pair electrons ~
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- bond pair electrons
i As discussed earlier in the bonding of water (H2
O) the
molecular shape and angle is described in the diagram below
Sul hur which is also an element from Grou 16 formed h dro en
sulphide H2S when sulphur react with hydrogen However unlike
water the bonding angles is much smaller compare to water This is
due to the difference of electronegativity and also the bond length
between O and S in molecule Since O is more electronegative
compare to S the O-H bond is pulled closer toward O Furthermorethe bond length of O-H is shorter compare to S-H As a result the
bonding pair-bonding pair electrons repulse greater with each other in
H-O-H and caused the angle become greater
ii Another example is between NH3 and PH3 The orbital
f
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diagram for both NH3 and PH3 are described below Both
nitrogen N and phosphorous P are from the same groupwhich is Group 15
Similar to the above case P in phosphine is less
electronegative than N in ammonia and bond length of N-H
is shorter than P-H As a result H is pull closer to N and
repulsion between bonding of H-N-H is greater compare toH-P-H and caused the angle between H-N-H is greater
compare to H-P-H
324 Valence Shell Electron Pair Repulsion (VSEPR) Theory
1 M l l t i th th di i l t f t
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1 Molecular geometry is the three-dimensional arrangement of atoms
in a molecule A moleculersquos geometry affects its physical andchemical properties such as melting point boiling point density
and the types of reactions it undergoes The basic concept of
VSEPR theory is based on the three general rules below
a) As far as electron-pair repulsion is concerned double bonds andtriple bonds can be treated like single bonds This approximation is
good for qualitative purposes However you should realize that in
ldquo rdquo
because there are two or three bonds between two atoms theelectron density occupies more space
b) If a molecule has two or more resonance structures we can apply
the VSEPR model to any one of them Formal charges are usually
not shown
c) The order of repulsion strength of lone pair and bond pair are
lone-pair amp lone-pair electrons repulsion are the strongest
followed by lone-pair amp bond-pair electrons repulsion while
bond-pair amp bond-pair electrons repulsion is the weakest
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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3 sp hybridisation ~ When two electron groups surround
the central atom we observe a linear shape which
means that the bonding orbitals must have a linear
orientation VB theory explains this by proposing thatmixing two nonequivalent orbitals of a central atom one
s and one p gives rise to two equivalent sp hybrid
orbitals that lie 1800 apart
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Beryllium which act as the
center atom has the valence
electron of 2s2
b) Excited state
One of the electron from 2s ispromote to 2p orbitals - equal
energy level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of a
linear shape
Molecular shape Linear
Angle between bond pair 1800
4 The concept of hybridisation is also useful to explain molecules with
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p y p
doubletriple bonds By using the concept of the direct overlapping
orbitals and side-touch lapping orbitals the formation of multiple
bonds in ethene C2H4 and ethyne C2H2 are described
a) Ethene C2H4 - Hybridisation take place for both carbon atoms in
ethene molecule is sp2 hybridisation
bull When both hybridised C is bonded together one of the hybridisedorbital overlapped directly between each other while the other two
hybridised orbitals overlapped directly with hydrogen atoms
z
atoms form a side-touch bond between each other and form anotherbond as shown in the diagram below
bull From the diagram with C=C there are two types of bond A sigma-
bond (σ-bond) is covalent bonds formed by orbitals overlapping end-
to-end with the electron density concentrated between the nuclei ofthe bonding atoms while the second type is called a pi bond (π-
bond) which is defined as a covalent bond formed by sideways
overlapping orbitals with electron density concentrated above and
below the plane of the nuclei of the bonding atoms
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital of 2s and two
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of atrigonal planar shape
Note that on unhybridised pz
orbital an electron is presence
Molecular shape Trigonal planar Angle between bond pair 1200
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b) Ethyne C H has the Lewis structure of H CequivC H which the
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b) Ethyne C2H2 has the Lewis structure of HminusCequivCminusH which the
bonding can be explain using sp hybridisation Table below
described how the hybridisation take place on each carbonatom and how the formation of triple bond occur
c) From the diagram formation of -CequivC- is due to the formation
of one sigma-bond by direct overlapping of one of the twohybrid orbitals on each C while the other two pi-bond bonds
are formed as a result of side-lapping of the each two
y z
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital from 2s and 2p
orbitals combined (hybrid) and
rearrange themselves to the
shape and orientation of a
linear shapeNote that on unhybridised py amp
pz orbitals electrons are
presence to form two
π-bonds Molecular shape Linear Angle between bond pair 1800
sp
pz py
5 Other examples of applications in valence bond theory includes the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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formation of nitrogen molecule N2 and hydrogen cyanide HCN
molecule
a) Nitrogen gas is Earths most abundant gas as it cover 78 of the
content of our air Nitrogen molecule is an inert gas thanks to its
short covalent bond and also its strong triple bond Therefore
a lot of heats are required to break the chemical bond of nitrogenbefore it can be applied in industries Using valence bond theory
the bonding of nitrogen molecule is explained in the diagram
b) So when the two hybridised nitrogen atom interacting amongeach other they hence formed a linear shape and two pi-bonds
(πminusbonds) are formed as a result of side-lapping of unhybridised
py and pz orbital respectively
N N
ExplanationEnergy level
diagramDiagram of orbitals
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diagram
a) Ground state
Both nitrogen whichact as the center atom
has the valence
electron of 2s22p3
b) Excited state One of the electron
from 2s is promote to 2p
orbitals - e ual ener
level
c) Hybridised state
One orbital of 2s and
one orbitals of 2p
combined (hybrid) and
rearrange themselves
to the shape and
orientation of a linear
shapeMolecular shape Linear
Angle between bond pair 1800
b) Whereas for hydrogen cyanide HCN both carbon atom and
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) y g y
nitrogen atom undergoes sp hybridisation in order to form a
linear structureExplanation
Energy level diagram for
carbon atom C
Energy level diagram for
nitrogen atom N
a) Ground state
Valence electronC 2s22p2
N 2s22p3
b) Excited state
One of the electron from
2s is promote to 2p
orbitals - equal energy
level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined
(hybrid) and form sp
hybridisation
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N
σσσσ
π
σσσσ
π
6 However there are a few limitation on valence bond theory such
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as when explaining the effect angle of bond-pair and bond-pair
electrons when there isare presence of lone pair electron in thecenter atom and also the difference of electronegativity
a) Bonding in ammonia NH3 and water H2O ~ Both NH3 and
H2O undergoes an arrangement similar to sp3 hybridisation
similar to that of C in methane Table below shows thehybridisation occur for nitrogen in ammonia and oxygen in water
Energy level diagram Energy level diagram
for nitrogen atom N for oxygen atom O
a) Ground state
Valence electron
N 2s22p3
O 2s2
2p4
b) Excited state
One of the electron
from 2s is promote to
2p orbitals
c) Hybridised state
One orbital of 2s and
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One orbital of 2s and
three orbitals from 2p
combined (hybrid) andrearrange themselves to
the shape and
orientation similar to that
of tetrahedral shape
Similar to arrangement in
tetrahedral
Similar to arrangement in
tetrahedral
of ammonia NH3 andwater H2O
Shape and angleShape trigonal pyramidal
Angle 1070
Shape bent
Angle 10450
Number of bond pair amp
lone pair electrons
Bond pair electron 3
Lone pair electron 1
Bond pair electron 2
Lone pair electron 2
bull The angle of bond pair - bond pair electrons in ammonia and0 0
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water are 1070 and 10450 respectively which is lesser than in
methane molecule (10950
) This can be explained by the factof the presence of lone-pair electrons in both ammonia and
water Since the lone pair - lone pair electron repulsion is
stronger than lone pair - bond pair electron repulsion than
bond pair - bond pair electron repulsion it is expected that therepulsion between lone pair - bond pair electrons in ammonia
is stronger than bond pair - bond pair electrons repulsion
hence caused the angle to squeeze to a smaller angle As
for water since there is a presence of lone pair - lone pair
electrons repulsion it results the bond pair - bond pair
electron repulsion to be much smaller hence caused the
angle to squeeze to a smaller angle Though valence bondtheory does not actually explain the hybridisation especially
when it relates to the repulsion occur involving lone pair
electron another theories shall be applied to study such
effect All these shall be discussed further in VSEPR theory
b) Effect of electronegativity and bonding angle between bond pair
- bond pair electrons ~
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- bond pair electrons
i As discussed earlier in the bonding of water (H2
O) the
molecular shape and angle is described in the diagram below
Sul hur which is also an element from Grou 16 formed h dro en
sulphide H2S when sulphur react with hydrogen However unlike
water the bonding angles is much smaller compare to water This is
due to the difference of electronegativity and also the bond length
between O and S in molecule Since O is more electronegative
compare to S the O-H bond is pulled closer toward O Furthermorethe bond length of O-H is shorter compare to S-H As a result the
bonding pair-bonding pair electrons repulse greater with each other in
H-O-H and caused the angle become greater
ii Another example is between NH3 and PH3 The orbital
f
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diagram for both NH3 and PH3 are described below Both
nitrogen N and phosphorous P are from the same groupwhich is Group 15
Similar to the above case P in phosphine is less
electronegative than N in ammonia and bond length of N-H
is shorter than P-H As a result H is pull closer to N and
repulsion between bonding of H-N-H is greater compare toH-P-H and caused the angle between H-N-H is greater
compare to H-P-H
324 Valence Shell Electron Pair Repulsion (VSEPR) Theory
1 M l l t i th th di i l t f t
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1 Molecular geometry is the three-dimensional arrangement of atoms
in a molecule A moleculersquos geometry affects its physical andchemical properties such as melting point boiling point density
and the types of reactions it undergoes The basic concept of
VSEPR theory is based on the three general rules below
a) As far as electron-pair repulsion is concerned double bonds andtriple bonds can be treated like single bonds This approximation is
good for qualitative purposes However you should realize that in
ldquo rdquo
because there are two or three bonds between two atoms theelectron density occupies more space
b) If a molecule has two or more resonance structures we can apply
the VSEPR model to any one of them Formal charges are usually
not shown
c) The order of repulsion strength of lone pair and bond pair are
lone-pair amp lone-pair electrons repulsion are the strongest
followed by lone-pair amp bond-pair electrons repulsion while
bond-pair amp bond-pair electrons repulsion is the weakest
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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3 sp hybridisation ~ When two electron groups surround
the central atom we observe a linear shape which
means that the bonding orbitals must have a linear
orientation VB theory explains this by proposing thatmixing two nonequivalent orbitals of a central atom one
s and one p gives rise to two equivalent sp hybrid
orbitals that lie 1800 apart
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Beryllium which act as the
center atom has the valence
electron of 2s2
b) Excited state
One of the electron from 2s ispromote to 2p orbitals - equal
energy level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of a
linear shape
Molecular shape Linear
Angle between bond pair 1800
4 The concept of hybridisation is also useful to explain molecules with
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p y p
doubletriple bonds By using the concept of the direct overlapping
orbitals and side-touch lapping orbitals the formation of multiple
bonds in ethene C2H4 and ethyne C2H2 are described
a) Ethene C2H4 - Hybridisation take place for both carbon atoms in
ethene molecule is sp2 hybridisation
bull When both hybridised C is bonded together one of the hybridisedorbital overlapped directly between each other while the other two
hybridised orbitals overlapped directly with hydrogen atoms
z
atoms form a side-touch bond between each other and form anotherbond as shown in the diagram below
bull From the diagram with C=C there are two types of bond A sigma-
bond (σ-bond) is covalent bonds formed by orbitals overlapping end-
to-end with the electron density concentrated between the nuclei ofthe bonding atoms while the second type is called a pi bond (π-
bond) which is defined as a covalent bond formed by sideways
overlapping orbitals with electron density concentrated above and
below the plane of the nuclei of the bonding atoms
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital of 2s and two
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of atrigonal planar shape
Note that on unhybridised pz
orbital an electron is presence
Molecular shape Trigonal planar Angle between bond pair 1200
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b) Ethyne C H has the Lewis structure of H CequivC H which the
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b) Ethyne C2H2 has the Lewis structure of HminusCequivCminusH which the
bonding can be explain using sp hybridisation Table below
described how the hybridisation take place on each carbonatom and how the formation of triple bond occur
c) From the diagram formation of -CequivC- is due to the formation
of one sigma-bond by direct overlapping of one of the twohybrid orbitals on each C while the other two pi-bond bonds
are formed as a result of side-lapping of the each two
y z
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital from 2s and 2p
orbitals combined (hybrid) and
rearrange themselves to the
shape and orientation of a
linear shapeNote that on unhybridised py amp
pz orbitals electrons are
presence to form two
π-bonds Molecular shape Linear Angle between bond pair 1800
sp
pz py
5 Other examples of applications in valence bond theory includes the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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formation of nitrogen molecule N2 and hydrogen cyanide HCN
molecule
a) Nitrogen gas is Earths most abundant gas as it cover 78 of the
content of our air Nitrogen molecule is an inert gas thanks to its
short covalent bond and also its strong triple bond Therefore
a lot of heats are required to break the chemical bond of nitrogenbefore it can be applied in industries Using valence bond theory
the bonding of nitrogen molecule is explained in the diagram
b) So when the two hybridised nitrogen atom interacting amongeach other they hence formed a linear shape and two pi-bonds
(πminusbonds) are formed as a result of side-lapping of unhybridised
py and pz orbital respectively
N N
ExplanationEnergy level
diagramDiagram of orbitals
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diagram
a) Ground state
Both nitrogen whichact as the center atom
has the valence
electron of 2s22p3
b) Excited state One of the electron
from 2s is promote to 2p
orbitals - e ual ener
level
c) Hybridised state
One orbital of 2s and
one orbitals of 2p
combined (hybrid) and
rearrange themselves
to the shape and
orientation of a linear
shapeMolecular shape Linear
Angle between bond pair 1800
b) Whereas for hydrogen cyanide HCN both carbon atom and
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) y g y
nitrogen atom undergoes sp hybridisation in order to form a
linear structureExplanation
Energy level diagram for
carbon atom C
Energy level diagram for
nitrogen atom N
a) Ground state
Valence electronC 2s22p2
N 2s22p3
b) Excited state
One of the electron from
2s is promote to 2p
orbitals - equal energy
level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined
(hybrid) and form sp
hybridisation
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N
σσσσ
π
σσσσ
π
6 However there are a few limitation on valence bond theory such
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as when explaining the effect angle of bond-pair and bond-pair
electrons when there isare presence of lone pair electron in thecenter atom and also the difference of electronegativity
a) Bonding in ammonia NH3 and water H2O ~ Both NH3 and
H2O undergoes an arrangement similar to sp3 hybridisation
similar to that of C in methane Table below shows thehybridisation occur for nitrogen in ammonia and oxygen in water
Energy level diagram Energy level diagram
for nitrogen atom N for oxygen atom O
a) Ground state
Valence electron
N 2s22p3
O 2s2
2p4
b) Excited state
One of the electron
from 2s is promote to
2p orbitals
c) Hybridised state
One orbital of 2s and
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One orbital of 2s and
three orbitals from 2p
combined (hybrid) andrearrange themselves to
the shape and
orientation similar to that
of tetrahedral shape
Similar to arrangement in
tetrahedral
Similar to arrangement in
tetrahedral
of ammonia NH3 andwater H2O
Shape and angleShape trigonal pyramidal
Angle 1070
Shape bent
Angle 10450
Number of bond pair amp
lone pair electrons
Bond pair electron 3
Lone pair electron 1
Bond pair electron 2
Lone pair electron 2
bull The angle of bond pair - bond pair electrons in ammonia and0 0
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water are 1070 and 10450 respectively which is lesser than in
methane molecule (10950
) This can be explained by the factof the presence of lone-pair electrons in both ammonia and
water Since the lone pair - lone pair electron repulsion is
stronger than lone pair - bond pair electron repulsion than
bond pair - bond pair electron repulsion it is expected that therepulsion between lone pair - bond pair electrons in ammonia
is stronger than bond pair - bond pair electrons repulsion
hence caused the angle to squeeze to a smaller angle As
for water since there is a presence of lone pair - lone pair
electrons repulsion it results the bond pair - bond pair
electron repulsion to be much smaller hence caused the
angle to squeeze to a smaller angle Though valence bondtheory does not actually explain the hybridisation especially
when it relates to the repulsion occur involving lone pair
electron another theories shall be applied to study such
effect All these shall be discussed further in VSEPR theory
b) Effect of electronegativity and bonding angle between bond pair
- bond pair electrons ~
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- bond pair electrons
i As discussed earlier in the bonding of water (H2
O) the
molecular shape and angle is described in the diagram below
Sul hur which is also an element from Grou 16 formed h dro en
sulphide H2S when sulphur react with hydrogen However unlike
water the bonding angles is much smaller compare to water This is
due to the difference of electronegativity and also the bond length
between O and S in molecule Since O is more electronegative
compare to S the O-H bond is pulled closer toward O Furthermorethe bond length of O-H is shorter compare to S-H As a result the
bonding pair-bonding pair electrons repulse greater with each other in
H-O-H and caused the angle become greater
ii Another example is between NH3 and PH3 The orbital
f
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diagram for both NH3 and PH3 are described below Both
nitrogen N and phosphorous P are from the same groupwhich is Group 15
Similar to the above case P in phosphine is less
electronegative than N in ammonia and bond length of N-H
is shorter than P-H As a result H is pull closer to N and
repulsion between bonding of H-N-H is greater compare toH-P-H and caused the angle between H-N-H is greater
compare to H-P-H
324 Valence Shell Electron Pair Repulsion (VSEPR) Theory
1 M l l t i th th di i l t f t
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1 Molecular geometry is the three-dimensional arrangement of atoms
in a molecule A moleculersquos geometry affects its physical andchemical properties such as melting point boiling point density
and the types of reactions it undergoes The basic concept of
VSEPR theory is based on the three general rules below
a) As far as electron-pair repulsion is concerned double bonds andtriple bonds can be treated like single bonds This approximation is
good for qualitative purposes However you should realize that in
ldquo rdquo
because there are two or three bonds between two atoms theelectron density occupies more space
b) If a molecule has two or more resonance structures we can apply
the VSEPR model to any one of them Formal charges are usually
not shown
c) The order of repulsion strength of lone pair and bond pair are
lone-pair amp lone-pair electrons repulsion are the strongest
followed by lone-pair amp bond-pair electrons repulsion while
bond-pair amp bond-pair electrons repulsion is the weakest
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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a) Ground state
Beryllium which act as the
center atom has the valence
electron of 2s2
b) Excited state
One of the electron from 2s ispromote to 2p orbitals - equal
energy level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of a
linear shape
Molecular shape Linear
Angle between bond pair 1800
4 The concept of hybridisation is also useful to explain molecules with
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p y p
doubletriple bonds By using the concept of the direct overlapping
orbitals and side-touch lapping orbitals the formation of multiple
bonds in ethene C2H4 and ethyne C2H2 are described
a) Ethene C2H4 - Hybridisation take place for both carbon atoms in
ethene molecule is sp2 hybridisation
bull When both hybridised C is bonded together one of the hybridisedorbital overlapped directly between each other while the other two
hybridised orbitals overlapped directly with hydrogen atoms
z
atoms form a side-touch bond between each other and form anotherbond as shown in the diagram below
bull From the diagram with C=C there are two types of bond A sigma-
bond (σ-bond) is covalent bonds formed by orbitals overlapping end-
to-end with the electron density concentrated between the nuclei ofthe bonding atoms while the second type is called a pi bond (π-
bond) which is defined as a covalent bond formed by sideways
overlapping orbitals with electron density concentrated above and
below the plane of the nuclei of the bonding atoms
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital of 2s and two
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of atrigonal planar shape
Note that on unhybridised pz
orbital an electron is presence
Molecular shape Trigonal planar Angle between bond pair 1200
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b) Ethyne C H has the Lewis structure of H CequivC H which the
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b) Ethyne C2H2 has the Lewis structure of HminusCequivCminusH which the
bonding can be explain using sp hybridisation Table below
described how the hybridisation take place on each carbonatom and how the formation of triple bond occur
c) From the diagram formation of -CequivC- is due to the formation
of one sigma-bond by direct overlapping of one of the twohybrid orbitals on each C while the other two pi-bond bonds
are formed as a result of side-lapping of the each two
y z
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital from 2s and 2p
orbitals combined (hybrid) and
rearrange themselves to the
shape and orientation of a
linear shapeNote that on unhybridised py amp
pz orbitals electrons are
presence to form two
π-bonds Molecular shape Linear Angle between bond pair 1800
sp
pz py
5 Other examples of applications in valence bond theory includes the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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formation of nitrogen molecule N2 and hydrogen cyanide HCN
molecule
a) Nitrogen gas is Earths most abundant gas as it cover 78 of the
content of our air Nitrogen molecule is an inert gas thanks to its
short covalent bond and also its strong triple bond Therefore
a lot of heats are required to break the chemical bond of nitrogenbefore it can be applied in industries Using valence bond theory
the bonding of nitrogen molecule is explained in the diagram
b) So when the two hybridised nitrogen atom interacting amongeach other they hence formed a linear shape and two pi-bonds
(πminusbonds) are formed as a result of side-lapping of unhybridised
py and pz orbital respectively
N N
ExplanationEnergy level
diagramDiagram of orbitals
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diagram
a) Ground state
Both nitrogen whichact as the center atom
has the valence
electron of 2s22p3
b) Excited state One of the electron
from 2s is promote to 2p
orbitals - e ual ener
level
c) Hybridised state
One orbital of 2s and
one orbitals of 2p
combined (hybrid) and
rearrange themselves
to the shape and
orientation of a linear
shapeMolecular shape Linear
Angle between bond pair 1800
b) Whereas for hydrogen cyanide HCN both carbon atom and
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) y g y
nitrogen atom undergoes sp hybridisation in order to form a
linear structureExplanation
Energy level diagram for
carbon atom C
Energy level diagram for
nitrogen atom N
a) Ground state
Valence electronC 2s22p2
N 2s22p3
b) Excited state
One of the electron from
2s is promote to 2p
orbitals - equal energy
level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined
(hybrid) and form sp
hybridisation
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N
σσσσ
π
σσσσ
π
6 However there are a few limitation on valence bond theory such
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as when explaining the effect angle of bond-pair and bond-pair
electrons when there isare presence of lone pair electron in thecenter atom and also the difference of electronegativity
a) Bonding in ammonia NH3 and water H2O ~ Both NH3 and
H2O undergoes an arrangement similar to sp3 hybridisation
similar to that of C in methane Table below shows thehybridisation occur for nitrogen in ammonia and oxygen in water
Energy level diagram Energy level diagram
for nitrogen atom N for oxygen atom O
a) Ground state
Valence electron
N 2s22p3
O 2s2
2p4
b) Excited state
One of the electron
from 2s is promote to
2p orbitals
c) Hybridised state
One orbital of 2s and
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One orbital of 2s and
three orbitals from 2p
combined (hybrid) andrearrange themselves to
the shape and
orientation similar to that
of tetrahedral shape
Similar to arrangement in
tetrahedral
Similar to arrangement in
tetrahedral
of ammonia NH3 andwater H2O
Shape and angleShape trigonal pyramidal
Angle 1070
Shape bent
Angle 10450
Number of bond pair amp
lone pair electrons
Bond pair electron 3
Lone pair electron 1
Bond pair electron 2
Lone pair electron 2
bull The angle of bond pair - bond pair electrons in ammonia and0 0
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water are 1070 and 10450 respectively which is lesser than in
methane molecule (10950
) This can be explained by the factof the presence of lone-pair electrons in both ammonia and
water Since the lone pair - lone pair electron repulsion is
stronger than lone pair - bond pair electron repulsion than
bond pair - bond pair electron repulsion it is expected that therepulsion between lone pair - bond pair electrons in ammonia
is stronger than bond pair - bond pair electrons repulsion
hence caused the angle to squeeze to a smaller angle As
for water since there is a presence of lone pair - lone pair
electrons repulsion it results the bond pair - bond pair
electron repulsion to be much smaller hence caused the
angle to squeeze to a smaller angle Though valence bondtheory does not actually explain the hybridisation especially
when it relates to the repulsion occur involving lone pair
electron another theories shall be applied to study such
effect All these shall be discussed further in VSEPR theory
b) Effect of electronegativity and bonding angle between bond pair
- bond pair electrons ~
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- bond pair electrons
i As discussed earlier in the bonding of water (H2
O) the
molecular shape and angle is described in the diagram below
Sul hur which is also an element from Grou 16 formed h dro en
sulphide H2S when sulphur react with hydrogen However unlike
water the bonding angles is much smaller compare to water This is
due to the difference of electronegativity and also the bond length
between O and S in molecule Since O is more electronegative
compare to S the O-H bond is pulled closer toward O Furthermorethe bond length of O-H is shorter compare to S-H As a result the
bonding pair-bonding pair electrons repulse greater with each other in
H-O-H and caused the angle become greater
ii Another example is between NH3 and PH3 The orbital
f
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diagram for both NH3 and PH3 are described below Both
nitrogen N and phosphorous P are from the same groupwhich is Group 15
Similar to the above case P in phosphine is less
electronegative than N in ammonia and bond length of N-H
is shorter than P-H As a result H is pull closer to N and
repulsion between bonding of H-N-H is greater compare toH-P-H and caused the angle between H-N-H is greater
compare to H-P-H
324 Valence Shell Electron Pair Repulsion (VSEPR) Theory
1 M l l t i th th di i l t f t
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1 Molecular geometry is the three-dimensional arrangement of atoms
in a molecule A moleculersquos geometry affects its physical andchemical properties such as melting point boiling point density
and the types of reactions it undergoes The basic concept of
VSEPR theory is based on the three general rules below
a) As far as electron-pair repulsion is concerned double bonds andtriple bonds can be treated like single bonds This approximation is
good for qualitative purposes However you should realize that in
ldquo rdquo
because there are two or three bonds between two atoms theelectron density occupies more space
b) If a molecule has two or more resonance structures we can apply
the VSEPR model to any one of them Formal charges are usually
not shown
c) The order of repulsion strength of lone pair and bond pair are
lone-pair amp lone-pair electrons repulsion are the strongest
followed by lone-pair amp bond-pair electrons repulsion while
bond-pair amp bond-pair electrons repulsion is the weakest
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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p y p
doubletriple bonds By using the concept of the direct overlapping
orbitals and side-touch lapping orbitals the formation of multiple
bonds in ethene C2H4 and ethyne C2H2 are described
a) Ethene C2H4 - Hybridisation take place for both carbon atoms in
ethene molecule is sp2 hybridisation
bull When both hybridised C is bonded together one of the hybridisedorbital overlapped directly between each other while the other two
hybridised orbitals overlapped directly with hydrogen atoms
z
atoms form a side-touch bond between each other and form anotherbond as shown in the diagram below
bull From the diagram with C=C there are two types of bond A sigma-
bond (σ-bond) is covalent bonds formed by orbitals overlapping end-
to-end with the electron density concentrated between the nuclei ofthe bonding atoms while the second type is called a pi bond (π-
bond) which is defined as a covalent bond formed by sideways
overlapping orbitals with electron density concentrated above and
below the plane of the nuclei of the bonding atoms
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital of 2s and two
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of atrigonal planar shape
Note that on unhybridised pz
orbital an electron is presence
Molecular shape Trigonal planar Angle between bond pair 1200
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b) Ethyne C H has the Lewis structure of H CequivC H which the
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b) Ethyne C2H2 has the Lewis structure of HminusCequivCminusH which the
bonding can be explain using sp hybridisation Table below
described how the hybridisation take place on each carbonatom and how the formation of triple bond occur
c) From the diagram formation of -CequivC- is due to the formation
of one sigma-bond by direct overlapping of one of the twohybrid orbitals on each C while the other two pi-bond bonds
are formed as a result of side-lapping of the each two
y z
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital from 2s and 2p
orbitals combined (hybrid) and
rearrange themselves to the
shape and orientation of a
linear shapeNote that on unhybridised py amp
pz orbitals electrons are
presence to form two
π-bonds Molecular shape Linear Angle between bond pair 1800
sp
pz py
5 Other examples of applications in valence bond theory includes the
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formation of nitrogen molecule N2 and hydrogen cyanide HCN
molecule
a) Nitrogen gas is Earths most abundant gas as it cover 78 of the
content of our air Nitrogen molecule is an inert gas thanks to its
short covalent bond and also its strong triple bond Therefore
a lot of heats are required to break the chemical bond of nitrogenbefore it can be applied in industries Using valence bond theory
the bonding of nitrogen molecule is explained in the diagram
b) So when the two hybridised nitrogen atom interacting amongeach other they hence formed a linear shape and two pi-bonds
(πminusbonds) are formed as a result of side-lapping of unhybridised
py and pz orbital respectively
N N
ExplanationEnergy level
diagramDiagram of orbitals
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diagram
a) Ground state
Both nitrogen whichact as the center atom
has the valence
electron of 2s22p3
b) Excited state One of the electron
from 2s is promote to 2p
orbitals - e ual ener
level
c) Hybridised state
One orbital of 2s and
one orbitals of 2p
combined (hybrid) and
rearrange themselves
to the shape and
orientation of a linear
shapeMolecular shape Linear
Angle between bond pair 1800
b) Whereas for hydrogen cyanide HCN both carbon atom and
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) y g y
nitrogen atom undergoes sp hybridisation in order to form a
linear structureExplanation
Energy level diagram for
carbon atom C
Energy level diagram for
nitrogen atom N
a) Ground state
Valence electronC 2s22p2
N 2s22p3
b) Excited state
One of the electron from
2s is promote to 2p
orbitals - equal energy
level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined
(hybrid) and form sp
hybridisation
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N
σσσσ
π
σσσσ
π
6 However there are a few limitation on valence bond theory such
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as when explaining the effect angle of bond-pair and bond-pair
electrons when there isare presence of lone pair electron in thecenter atom and also the difference of electronegativity
a) Bonding in ammonia NH3 and water H2O ~ Both NH3 and
H2O undergoes an arrangement similar to sp3 hybridisation
similar to that of C in methane Table below shows thehybridisation occur for nitrogen in ammonia and oxygen in water
Energy level diagram Energy level diagram
for nitrogen atom N for oxygen atom O
a) Ground state
Valence electron
N 2s22p3
O 2s2
2p4
b) Excited state
One of the electron
from 2s is promote to
2p orbitals
c) Hybridised state
One orbital of 2s and
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One orbital of 2s and
three orbitals from 2p
combined (hybrid) andrearrange themselves to
the shape and
orientation similar to that
of tetrahedral shape
Similar to arrangement in
tetrahedral
Similar to arrangement in
tetrahedral
of ammonia NH3 andwater H2O
Shape and angleShape trigonal pyramidal
Angle 1070
Shape bent
Angle 10450
Number of bond pair amp
lone pair electrons
Bond pair electron 3
Lone pair electron 1
Bond pair electron 2
Lone pair electron 2
bull The angle of bond pair - bond pair electrons in ammonia and0 0
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water are 1070 and 10450 respectively which is lesser than in
methane molecule (10950
) This can be explained by the factof the presence of lone-pair electrons in both ammonia and
water Since the lone pair - lone pair electron repulsion is
stronger than lone pair - bond pair electron repulsion than
bond pair - bond pair electron repulsion it is expected that therepulsion between lone pair - bond pair electrons in ammonia
is stronger than bond pair - bond pair electrons repulsion
hence caused the angle to squeeze to a smaller angle As
for water since there is a presence of lone pair - lone pair
electrons repulsion it results the bond pair - bond pair
electron repulsion to be much smaller hence caused the
angle to squeeze to a smaller angle Though valence bondtheory does not actually explain the hybridisation especially
when it relates to the repulsion occur involving lone pair
electron another theories shall be applied to study such
effect All these shall be discussed further in VSEPR theory
b) Effect of electronegativity and bonding angle between bond pair
- bond pair electrons ~
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- bond pair electrons
i As discussed earlier in the bonding of water (H2
O) the
molecular shape and angle is described in the diagram below
Sul hur which is also an element from Grou 16 formed h dro en
sulphide H2S when sulphur react with hydrogen However unlike
water the bonding angles is much smaller compare to water This is
due to the difference of electronegativity and also the bond length
between O and S in molecule Since O is more electronegative
compare to S the O-H bond is pulled closer toward O Furthermorethe bond length of O-H is shorter compare to S-H As a result the
bonding pair-bonding pair electrons repulse greater with each other in
H-O-H and caused the angle become greater
ii Another example is between NH3 and PH3 The orbital
f
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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diagram for both NH3 and PH3 are described below Both
nitrogen N and phosphorous P are from the same groupwhich is Group 15
Similar to the above case P in phosphine is less
electronegative than N in ammonia and bond length of N-H
is shorter than P-H As a result H is pull closer to N and
repulsion between bonding of H-N-H is greater compare toH-P-H and caused the angle between H-N-H is greater
compare to H-P-H
324 Valence Shell Electron Pair Repulsion (VSEPR) Theory
1 M l l t i th th di i l t f t
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1 Molecular geometry is the three-dimensional arrangement of atoms
in a molecule A moleculersquos geometry affects its physical andchemical properties such as melting point boiling point density
and the types of reactions it undergoes The basic concept of
VSEPR theory is based on the three general rules below
a) As far as electron-pair repulsion is concerned double bonds andtriple bonds can be treated like single bonds This approximation is
good for qualitative purposes However you should realize that in
ldquo rdquo
because there are two or three bonds between two atoms theelectron density occupies more space
b) If a molecule has two or more resonance structures we can apply
the VSEPR model to any one of them Formal charges are usually
not shown
c) The order of repulsion strength of lone pair and bond pair are
lone-pair amp lone-pair electrons repulsion are the strongest
followed by lone-pair amp bond-pair electrons repulsion while
bond-pair amp bond-pair electrons repulsion is the weakest
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital of 2s and two
orbitals of 2p combined (hybrid)
and rearrange themselves to
the shape and orientation of atrigonal planar shape
Note that on unhybridised pz
orbital an electron is presence
Molecular shape Trigonal planar Angle between bond pair 1200
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b) Ethyne C H has the Lewis structure of H CequivC H which the
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b) Ethyne C2H2 has the Lewis structure of HminusCequivCminusH which the
bonding can be explain using sp hybridisation Table below
described how the hybridisation take place on each carbonatom and how the formation of triple bond occur
c) From the diagram formation of -CequivC- is due to the formation
of one sigma-bond by direct overlapping of one of the twohybrid orbitals on each C while the other two pi-bond bonds
are formed as a result of side-lapping of the each two
y z
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital from 2s and 2p
orbitals combined (hybrid) and
rearrange themselves to the
shape and orientation of a
linear shapeNote that on unhybridised py amp
pz orbitals electrons are
presence to form two
π-bonds Molecular shape Linear Angle between bond pair 1800
sp
pz py
5 Other examples of applications in valence bond theory includes the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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formation of nitrogen molecule N2 and hydrogen cyanide HCN
molecule
a) Nitrogen gas is Earths most abundant gas as it cover 78 of the
content of our air Nitrogen molecule is an inert gas thanks to its
short covalent bond and also its strong triple bond Therefore
a lot of heats are required to break the chemical bond of nitrogenbefore it can be applied in industries Using valence bond theory
the bonding of nitrogen molecule is explained in the diagram
b) So when the two hybridised nitrogen atom interacting amongeach other they hence formed a linear shape and two pi-bonds
(πminusbonds) are formed as a result of side-lapping of unhybridised
py and pz orbital respectively
N N
ExplanationEnergy level
diagramDiagram of orbitals
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diagram
a) Ground state
Both nitrogen whichact as the center atom
has the valence
electron of 2s22p3
b) Excited state One of the electron
from 2s is promote to 2p
orbitals - e ual ener
level
c) Hybridised state
One orbital of 2s and
one orbitals of 2p
combined (hybrid) and
rearrange themselves
to the shape and
orientation of a linear
shapeMolecular shape Linear
Angle between bond pair 1800
b) Whereas for hydrogen cyanide HCN both carbon atom and
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) y g y
nitrogen atom undergoes sp hybridisation in order to form a
linear structureExplanation
Energy level diagram for
carbon atom C
Energy level diagram for
nitrogen atom N
a) Ground state
Valence electronC 2s22p2
N 2s22p3
b) Excited state
One of the electron from
2s is promote to 2p
orbitals - equal energy
level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined
(hybrid) and form sp
hybridisation
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N
σσσσ
π
σσσσ
π
6 However there are a few limitation on valence bond theory such
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as when explaining the effect angle of bond-pair and bond-pair
electrons when there isare presence of lone pair electron in thecenter atom and also the difference of electronegativity
a) Bonding in ammonia NH3 and water H2O ~ Both NH3 and
H2O undergoes an arrangement similar to sp3 hybridisation
similar to that of C in methane Table below shows thehybridisation occur for nitrogen in ammonia and oxygen in water
Energy level diagram Energy level diagram
for nitrogen atom N for oxygen atom O
a) Ground state
Valence electron
N 2s22p3
O 2s2
2p4
b) Excited state
One of the electron
from 2s is promote to
2p orbitals
c) Hybridised state
One orbital of 2s and
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One orbital of 2s and
three orbitals from 2p
combined (hybrid) andrearrange themselves to
the shape and
orientation similar to that
of tetrahedral shape
Similar to arrangement in
tetrahedral
Similar to arrangement in
tetrahedral
of ammonia NH3 andwater H2O
Shape and angleShape trigonal pyramidal
Angle 1070
Shape bent
Angle 10450
Number of bond pair amp
lone pair electrons
Bond pair electron 3
Lone pair electron 1
Bond pair electron 2
Lone pair electron 2
bull The angle of bond pair - bond pair electrons in ammonia and0 0
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water are 1070 and 10450 respectively which is lesser than in
methane molecule (10950
) This can be explained by the factof the presence of lone-pair electrons in both ammonia and
water Since the lone pair - lone pair electron repulsion is
stronger than lone pair - bond pair electron repulsion than
bond pair - bond pair electron repulsion it is expected that therepulsion between lone pair - bond pair electrons in ammonia
is stronger than bond pair - bond pair electrons repulsion
hence caused the angle to squeeze to a smaller angle As
for water since there is a presence of lone pair - lone pair
electrons repulsion it results the bond pair - bond pair
electron repulsion to be much smaller hence caused the
angle to squeeze to a smaller angle Though valence bondtheory does not actually explain the hybridisation especially
when it relates to the repulsion occur involving lone pair
electron another theories shall be applied to study such
effect All these shall be discussed further in VSEPR theory
b) Effect of electronegativity and bonding angle between bond pair
- bond pair electrons ~
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- bond pair electrons
i As discussed earlier in the bonding of water (H2
O) the
molecular shape and angle is described in the diagram below
Sul hur which is also an element from Grou 16 formed h dro en
sulphide H2S when sulphur react with hydrogen However unlike
water the bonding angles is much smaller compare to water This is
due to the difference of electronegativity and also the bond length
between O and S in molecule Since O is more electronegative
compare to S the O-H bond is pulled closer toward O Furthermorethe bond length of O-H is shorter compare to S-H As a result the
bonding pair-bonding pair electrons repulse greater with each other in
H-O-H and caused the angle become greater
ii Another example is between NH3 and PH3 The orbital
f
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diagram for both NH3 and PH3 are described below Both
nitrogen N and phosphorous P are from the same groupwhich is Group 15
Similar to the above case P in phosphine is less
electronegative than N in ammonia and bond length of N-H
is shorter than P-H As a result H is pull closer to N and
repulsion between bonding of H-N-H is greater compare toH-P-H and caused the angle between H-N-H is greater
compare to H-P-H
324 Valence Shell Electron Pair Repulsion (VSEPR) Theory
1 M l l t i th th di i l t f t
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1 Molecular geometry is the three-dimensional arrangement of atoms
in a molecule A moleculersquos geometry affects its physical andchemical properties such as melting point boiling point density
and the types of reactions it undergoes The basic concept of
VSEPR theory is based on the three general rules below
a) As far as electron-pair repulsion is concerned double bonds andtriple bonds can be treated like single bonds This approximation is
good for qualitative purposes However you should realize that in
ldquo rdquo
because there are two or three bonds between two atoms theelectron density occupies more space
b) If a molecule has two or more resonance structures we can apply
the VSEPR model to any one of them Formal charges are usually
not shown
c) The order of repulsion strength of lone pair and bond pair are
lone-pair amp lone-pair electrons repulsion are the strongest
followed by lone-pair amp bond-pair electrons repulsion while
bond-pair amp bond-pair electrons repulsion is the weakest
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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b) Ethyne C H has the Lewis structure of H CequivC H which the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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b) Ethyne C2H2 has the Lewis structure of HminusCequivCminusH which the
bonding can be explain using sp hybridisation Table below
described how the hybridisation take place on each carbonatom and how the formation of triple bond occur
c) From the diagram formation of -CequivC- is due to the formation
of one sigma-bond by direct overlapping of one of the twohybrid orbitals on each C while the other two pi-bond bonds
are formed as a result of side-lapping of the each two
y z
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital from 2s and 2p
orbitals combined (hybrid) and
rearrange themselves to the
shape and orientation of a
linear shapeNote that on unhybridised py amp
pz orbitals electrons are
presence to form two
π-bonds Molecular shape Linear Angle between bond pair 1800
sp
pz py
5 Other examples of applications in valence bond theory includes the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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formation of nitrogen molecule N2 and hydrogen cyanide HCN
molecule
a) Nitrogen gas is Earths most abundant gas as it cover 78 of the
content of our air Nitrogen molecule is an inert gas thanks to its
short covalent bond and also its strong triple bond Therefore
a lot of heats are required to break the chemical bond of nitrogenbefore it can be applied in industries Using valence bond theory
the bonding of nitrogen molecule is explained in the diagram
b) So when the two hybridised nitrogen atom interacting amongeach other they hence formed a linear shape and two pi-bonds
(πminusbonds) are formed as a result of side-lapping of unhybridised
py and pz orbital respectively
N N
ExplanationEnergy level
diagramDiagram of orbitals
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diagram
a) Ground state
Both nitrogen whichact as the center atom
has the valence
electron of 2s22p3
b) Excited state One of the electron
from 2s is promote to 2p
orbitals - e ual ener
level
c) Hybridised state
One orbital of 2s and
one orbitals of 2p
combined (hybrid) and
rearrange themselves
to the shape and
orientation of a linear
shapeMolecular shape Linear
Angle between bond pair 1800
b) Whereas for hydrogen cyanide HCN both carbon atom and
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) y g y
nitrogen atom undergoes sp hybridisation in order to form a
linear structureExplanation
Energy level diagram for
carbon atom C
Energy level diagram for
nitrogen atom N
a) Ground state
Valence electronC 2s22p2
N 2s22p3
b) Excited state
One of the electron from
2s is promote to 2p
orbitals - equal energy
level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined
(hybrid) and form sp
hybridisation
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N
σσσσ
π
σσσσ
π
6 However there are a few limitation on valence bond theory such
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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as when explaining the effect angle of bond-pair and bond-pair
electrons when there isare presence of lone pair electron in thecenter atom and also the difference of electronegativity
a) Bonding in ammonia NH3 and water H2O ~ Both NH3 and
H2O undergoes an arrangement similar to sp3 hybridisation
similar to that of C in methane Table below shows thehybridisation occur for nitrogen in ammonia and oxygen in water
Energy level diagram Energy level diagram
for nitrogen atom N for oxygen atom O
a) Ground state
Valence electron
N 2s22p3
O 2s2
2p4
b) Excited state
One of the electron
from 2s is promote to
2p orbitals
c) Hybridised state
One orbital of 2s and
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One orbital of 2s and
three orbitals from 2p
combined (hybrid) andrearrange themselves to
the shape and
orientation similar to that
of tetrahedral shape
Similar to arrangement in
tetrahedral
Similar to arrangement in
tetrahedral
of ammonia NH3 andwater H2O
Shape and angleShape trigonal pyramidal
Angle 1070
Shape bent
Angle 10450
Number of bond pair amp
lone pair electrons
Bond pair electron 3
Lone pair electron 1
Bond pair electron 2
Lone pair electron 2
bull The angle of bond pair - bond pair electrons in ammonia and0 0
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water are 1070 and 10450 respectively which is lesser than in
methane molecule (10950
) This can be explained by the factof the presence of lone-pair electrons in both ammonia and
water Since the lone pair - lone pair electron repulsion is
stronger than lone pair - bond pair electron repulsion than
bond pair - bond pair electron repulsion it is expected that therepulsion between lone pair - bond pair electrons in ammonia
is stronger than bond pair - bond pair electrons repulsion
hence caused the angle to squeeze to a smaller angle As
for water since there is a presence of lone pair - lone pair
electrons repulsion it results the bond pair - bond pair
electron repulsion to be much smaller hence caused the
angle to squeeze to a smaller angle Though valence bondtheory does not actually explain the hybridisation especially
when it relates to the repulsion occur involving lone pair
electron another theories shall be applied to study such
effect All these shall be discussed further in VSEPR theory
b) Effect of electronegativity and bonding angle between bond pair
- bond pair electrons ~
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- bond pair electrons
i As discussed earlier in the bonding of water (H2
O) the
molecular shape and angle is described in the diagram below
Sul hur which is also an element from Grou 16 formed h dro en
sulphide H2S when sulphur react with hydrogen However unlike
water the bonding angles is much smaller compare to water This is
due to the difference of electronegativity and also the bond length
between O and S in molecule Since O is more electronegative
compare to S the O-H bond is pulled closer toward O Furthermorethe bond length of O-H is shorter compare to S-H As a result the
bonding pair-bonding pair electrons repulse greater with each other in
H-O-H and caused the angle become greater
ii Another example is between NH3 and PH3 The orbital
f
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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diagram for both NH3 and PH3 are described below Both
nitrogen N and phosphorous P are from the same groupwhich is Group 15
Similar to the above case P in phosphine is less
electronegative than N in ammonia and bond length of N-H
is shorter than P-H As a result H is pull closer to N and
repulsion between bonding of H-N-H is greater compare toH-P-H and caused the angle between H-N-H is greater
compare to H-P-H
324 Valence Shell Electron Pair Repulsion (VSEPR) Theory
1 M l l t i th th di i l t f t
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1 Molecular geometry is the three-dimensional arrangement of atoms
in a molecule A moleculersquos geometry affects its physical andchemical properties such as melting point boiling point density
and the types of reactions it undergoes The basic concept of
VSEPR theory is based on the three general rules below
a) As far as electron-pair repulsion is concerned double bonds andtriple bonds can be treated like single bonds This approximation is
good for qualitative purposes However you should realize that in
ldquo rdquo
because there are two or three bonds between two atoms theelectron density occupies more space
b) If a molecule has two or more resonance structures we can apply
the VSEPR model to any one of them Formal charges are usually
not shown
c) The order of repulsion strength of lone pair and bond pair are
lone-pair amp lone-pair electrons repulsion are the strongest
followed by lone-pair amp bond-pair electrons repulsion while
bond-pair amp bond-pair electrons repulsion is the weakest
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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b) Ethyne C2H2 has the Lewis structure of HminusCequivCminusH which the
bonding can be explain using sp hybridisation Table below
described how the hybridisation take place on each carbonatom and how the formation of triple bond occur
c) From the diagram formation of -CequivC- is due to the formation
of one sigma-bond by direct overlapping of one of the twohybrid orbitals on each C while the other two pi-bond bonds
are formed as a result of side-lapping of the each two
y z
Explanation Energy level diagram Diagram of orbitals
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital from 2s and 2p
orbitals combined (hybrid) and
rearrange themselves to the
shape and orientation of a
linear shapeNote that on unhybridised py amp
pz orbitals electrons are
presence to form two
π-bonds Molecular shape Linear Angle between bond pair 1800
sp
pz py
5 Other examples of applications in valence bond theory includes the
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formation of nitrogen molecule N2 and hydrogen cyanide HCN
molecule
a) Nitrogen gas is Earths most abundant gas as it cover 78 of the
content of our air Nitrogen molecule is an inert gas thanks to its
short covalent bond and also its strong triple bond Therefore
a lot of heats are required to break the chemical bond of nitrogenbefore it can be applied in industries Using valence bond theory
the bonding of nitrogen molecule is explained in the diagram
b) So when the two hybridised nitrogen atom interacting amongeach other they hence formed a linear shape and two pi-bonds
(πminusbonds) are formed as a result of side-lapping of unhybridised
py and pz orbital respectively
N N
ExplanationEnergy level
diagramDiagram of orbitals
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diagram
a) Ground state
Both nitrogen whichact as the center atom
has the valence
electron of 2s22p3
b) Excited state One of the electron
from 2s is promote to 2p
orbitals - e ual ener
level
c) Hybridised state
One orbital of 2s and
one orbitals of 2p
combined (hybrid) and
rearrange themselves
to the shape and
orientation of a linear
shapeMolecular shape Linear
Angle between bond pair 1800
b) Whereas for hydrogen cyanide HCN both carbon atom and
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) y g y
nitrogen atom undergoes sp hybridisation in order to form a
linear structureExplanation
Energy level diagram for
carbon atom C
Energy level diagram for
nitrogen atom N
a) Ground state
Valence electronC 2s22p2
N 2s22p3
b) Excited state
One of the electron from
2s is promote to 2p
orbitals - equal energy
level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined
(hybrid) and form sp
hybridisation
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N
σσσσ
π
σσσσ
π
6 However there are a few limitation on valence bond theory such
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as when explaining the effect angle of bond-pair and bond-pair
electrons when there isare presence of lone pair electron in thecenter atom and also the difference of electronegativity
a) Bonding in ammonia NH3 and water H2O ~ Both NH3 and
H2O undergoes an arrangement similar to sp3 hybridisation
similar to that of C in methane Table below shows thehybridisation occur for nitrogen in ammonia and oxygen in water
Energy level diagram Energy level diagram
for nitrogen atom N for oxygen atom O
a) Ground state
Valence electron
N 2s22p3
O 2s2
2p4
b) Excited state
One of the electron
from 2s is promote to
2p orbitals
c) Hybridised state
One orbital of 2s and
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One orbital of 2s and
three orbitals from 2p
combined (hybrid) andrearrange themselves to
the shape and
orientation similar to that
of tetrahedral shape
Similar to arrangement in
tetrahedral
Similar to arrangement in
tetrahedral
of ammonia NH3 andwater H2O
Shape and angleShape trigonal pyramidal
Angle 1070
Shape bent
Angle 10450
Number of bond pair amp
lone pair electrons
Bond pair electron 3
Lone pair electron 1
Bond pair electron 2
Lone pair electron 2
bull The angle of bond pair - bond pair electrons in ammonia and0 0
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water are 1070 and 10450 respectively which is lesser than in
methane molecule (10950
) This can be explained by the factof the presence of lone-pair electrons in both ammonia and
water Since the lone pair - lone pair electron repulsion is
stronger than lone pair - bond pair electron repulsion than
bond pair - bond pair electron repulsion it is expected that therepulsion between lone pair - bond pair electrons in ammonia
is stronger than bond pair - bond pair electrons repulsion
hence caused the angle to squeeze to a smaller angle As
for water since there is a presence of lone pair - lone pair
electrons repulsion it results the bond pair - bond pair
electron repulsion to be much smaller hence caused the
angle to squeeze to a smaller angle Though valence bondtheory does not actually explain the hybridisation especially
when it relates to the repulsion occur involving lone pair
electron another theories shall be applied to study such
effect All these shall be discussed further in VSEPR theory
b) Effect of electronegativity and bonding angle between bond pair
- bond pair electrons ~
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- bond pair electrons
i As discussed earlier in the bonding of water (H2
O) the
molecular shape and angle is described in the diagram below
Sul hur which is also an element from Grou 16 formed h dro en
sulphide H2S when sulphur react with hydrogen However unlike
water the bonding angles is much smaller compare to water This is
due to the difference of electronegativity and also the bond length
between O and S in molecule Since O is more electronegative
compare to S the O-H bond is pulled closer toward O Furthermorethe bond length of O-H is shorter compare to S-H As a result the
bonding pair-bonding pair electrons repulse greater with each other in
H-O-H and caused the angle become greater
ii Another example is between NH3 and PH3 The orbital
f
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diagram for both NH3 and PH3 are described below Both
nitrogen N and phosphorous P are from the same groupwhich is Group 15
Similar to the above case P in phosphine is less
electronegative than N in ammonia and bond length of N-H
is shorter than P-H As a result H is pull closer to N and
repulsion between bonding of H-N-H is greater compare toH-P-H and caused the angle between H-N-H is greater
compare to H-P-H
324 Valence Shell Electron Pair Repulsion (VSEPR) Theory
1 M l l t i th th di i l t f t
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1 Molecular geometry is the three-dimensional arrangement of atoms
in a molecule A moleculersquos geometry affects its physical andchemical properties such as melting point boiling point density
and the types of reactions it undergoes The basic concept of
VSEPR theory is based on the three general rules below
a) As far as electron-pair repulsion is concerned double bonds andtriple bonds can be treated like single bonds This approximation is
good for qualitative purposes However you should realize that in
ldquo rdquo
because there are two or three bonds between two atoms theelectron density occupies more space
b) If a molecule has two or more resonance structures we can apply
the VSEPR model to any one of them Formal charges are usually
not shown
c) The order of repulsion strength of lone pair and bond pair are
lone-pair amp lone-pair electrons repulsion are the strongest
followed by lone-pair amp bond-pair electrons repulsion while
bond-pair amp bond-pair electrons repulsion is the weakest
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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a) Ground state
Carbon which act as the center
atom has the valence electron
of 2s22p2
b) Excited state
One of the electron from 2s is
promote to 2p orbitals - equalenergy level
c) Hybridised state
One orbital from 2s and 2p
orbitals combined (hybrid) and
rearrange themselves to the
shape and orientation of a
linear shapeNote that on unhybridised py amp
pz orbitals electrons are
presence to form two
π-bonds Molecular shape Linear Angle between bond pair 1800
sp
pz py
5 Other examples of applications in valence bond theory includes the
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formation of nitrogen molecule N2 and hydrogen cyanide HCN
molecule
a) Nitrogen gas is Earths most abundant gas as it cover 78 of the
content of our air Nitrogen molecule is an inert gas thanks to its
short covalent bond and also its strong triple bond Therefore
a lot of heats are required to break the chemical bond of nitrogenbefore it can be applied in industries Using valence bond theory
the bonding of nitrogen molecule is explained in the diagram
b) So when the two hybridised nitrogen atom interacting amongeach other they hence formed a linear shape and two pi-bonds
(πminusbonds) are formed as a result of side-lapping of unhybridised
py and pz orbital respectively
N N
ExplanationEnergy level
diagramDiagram of orbitals
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diagram
a) Ground state
Both nitrogen whichact as the center atom
has the valence
electron of 2s22p3
b) Excited state One of the electron
from 2s is promote to 2p
orbitals - e ual ener
level
c) Hybridised state
One orbital of 2s and
one orbitals of 2p
combined (hybrid) and
rearrange themselves
to the shape and
orientation of a linear
shapeMolecular shape Linear
Angle between bond pair 1800
b) Whereas for hydrogen cyanide HCN both carbon atom and
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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) y g y
nitrogen atom undergoes sp hybridisation in order to form a
linear structureExplanation
Energy level diagram for
carbon atom C
Energy level diagram for
nitrogen atom N
a) Ground state
Valence electronC 2s22p2
N 2s22p3
b) Excited state
One of the electron from
2s is promote to 2p
orbitals - equal energy
level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined
(hybrid) and form sp
hybridisation
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N
σσσσ
π
σσσσ
π
6 However there are a few limitation on valence bond theory such
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as when explaining the effect angle of bond-pair and bond-pair
electrons when there isare presence of lone pair electron in thecenter atom and also the difference of electronegativity
a) Bonding in ammonia NH3 and water H2O ~ Both NH3 and
H2O undergoes an arrangement similar to sp3 hybridisation
similar to that of C in methane Table below shows thehybridisation occur for nitrogen in ammonia and oxygen in water
Energy level diagram Energy level diagram
for nitrogen atom N for oxygen atom O
a) Ground state
Valence electron
N 2s22p3
O 2s2
2p4
b) Excited state
One of the electron
from 2s is promote to
2p orbitals
c) Hybridised state
One orbital of 2s and
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One orbital of 2s and
three orbitals from 2p
combined (hybrid) andrearrange themselves to
the shape and
orientation similar to that
of tetrahedral shape
Similar to arrangement in
tetrahedral
Similar to arrangement in
tetrahedral
of ammonia NH3 andwater H2O
Shape and angleShape trigonal pyramidal
Angle 1070
Shape bent
Angle 10450
Number of bond pair amp
lone pair electrons
Bond pair electron 3
Lone pair electron 1
Bond pair electron 2
Lone pair electron 2
bull The angle of bond pair - bond pair electrons in ammonia and0 0
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water are 1070 and 10450 respectively which is lesser than in
methane molecule (10950
) This can be explained by the factof the presence of lone-pair electrons in both ammonia and
water Since the lone pair - lone pair electron repulsion is
stronger than lone pair - bond pair electron repulsion than
bond pair - bond pair electron repulsion it is expected that therepulsion between lone pair - bond pair electrons in ammonia
is stronger than bond pair - bond pair electrons repulsion
hence caused the angle to squeeze to a smaller angle As
for water since there is a presence of lone pair - lone pair
electrons repulsion it results the bond pair - bond pair
electron repulsion to be much smaller hence caused the
angle to squeeze to a smaller angle Though valence bondtheory does not actually explain the hybridisation especially
when it relates to the repulsion occur involving lone pair
electron another theories shall be applied to study such
effect All these shall be discussed further in VSEPR theory
b) Effect of electronegativity and bonding angle between bond pair
- bond pair electrons ~
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- bond pair electrons
i As discussed earlier in the bonding of water (H2
O) the
molecular shape and angle is described in the diagram below
Sul hur which is also an element from Grou 16 formed h dro en
sulphide H2S when sulphur react with hydrogen However unlike
water the bonding angles is much smaller compare to water This is
due to the difference of electronegativity and also the bond length
between O and S in molecule Since O is more electronegative
compare to S the O-H bond is pulled closer toward O Furthermorethe bond length of O-H is shorter compare to S-H As a result the
bonding pair-bonding pair electrons repulse greater with each other in
H-O-H and caused the angle become greater
ii Another example is between NH3 and PH3 The orbital
f
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diagram for both NH3 and PH3 are described below Both
nitrogen N and phosphorous P are from the same groupwhich is Group 15
Similar to the above case P in phosphine is less
electronegative than N in ammonia and bond length of N-H
is shorter than P-H As a result H is pull closer to N and
repulsion between bonding of H-N-H is greater compare toH-P-H and caused the angle between H-N-H is greater
compare to H-P-H
324 Valence Shell Electron Pair Repulsion (VSEPR) Theory
1 M l l t i th th di i l t f t
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1 Molecular geometry is the three-dimensional arrangement of atoms
in a molecule A moleculersquos geometry affects its physical andchemical properties such as melting point boiling point density
and the types of reactions it undergoes The basic concept of
VSEPR theory is based on the three general rules below
a) As far as electron-pair repulsion is concerned double bonds andtriple bonds can be treated like single bonds This approximation is
good for qualitative purposes However you should realize that in
ldquo rdquo
because there are two or three bonds between two atoms theelectron density occupies more space
b) If a molecule has two or more resonance structures we can apply
the VSEPR model to any one of them Formal charges are usually
not shown
c) The order of repulsion strength of lone pair and bond pair are
lone-pair amp lone-pair electrons repulsion are the strongest
followed by lone-pair amp bond-pair electrons repulsion while
bond-pair amp bond-pair electrons repulsion is the weakest
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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formation of nitrogen molecule N2 and hydrogen cyanide HCN
molecule
a) Nitrogen gas is Earths most abundant gas as it cover 78 of the
content of our air Nitrogen molecule is an inert gas thanks to its
short covalent bond and also its strong triple bond Therefore
a lot of heats are required to break the chemical bond of nitrogenbefore it can be applied in industries Using valence bond theory
the bonding of nitrogen molecule is explained in the diagram
b) So when the two hybridised nitrogen atom interacting amongeach other they hence formed a linear shape and two pi-bonds
(πminusbonds) are formed as a result of side-lapping of unhybridised
py and pz orbital respectively
N N
ExplanationEnergy level
diagramDiagram of orbitals
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diagram
a) Ground state
Both nitrogen whichact as the center atom
has the valence
electron of 2s22p3
b) Excited state One of the electron
from 2s is promote to 2p
orbitals - e ual ener
level
c) Hybridised state
One orbital of 2s and
one orbitals of 2p
combined (hybrid) and
rearrange themselves
to the shape and
orientation of a linear
shapeMolecular shape Linear
Angle between bond pair 1800
b) Whereas for hydrogen cyanide HCN both carbon atom and
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) y g y
nitrogen atom undergoes sp hybridisation in order to form a
linear structureExplanation
Energy level diagram for
carbon atom C
Energy level diagram for
nitrogen atom N
a) Ground state
Valence electronC 2s22p2
N 2s22p3
b) Excited state
One of the electron from
2s is promote to 2p
orbitals - equal energy
level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined
(hybrid) and form sp
hybridisation
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N
σσσσ
π
σσσσ
π
6 However there are a few limitation on valence bond theory such
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as when explaining the effect angle of bond-pair and bond-pair
electrons when there isare presence of lone pair electron in thecenter atom and also the difference of electronegativity
a) Bonding in ammonia NH3 and water H2O ~ Both NH3 and
H2O undergoes an arrangement similar to sp3 hybridisation
similar to that of C in methane Table below shows thehybridisation occur for nitrogen in ammonia and oxygen in water
Energy level diagram Energy level diagram
for nitrogen atom N for oxygen atom O
a) Ground state
Valence electron
N 2s22p3
O 2s2
2p4
b) Excited state
One of the electron
from 2s is promote to
2p orbitals
c) Hybridised state
One orbital of 2s and
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One orbital of 2s and
three orbitals from 2p
combined (hybrid) andrearrange themselves to
the shape and
orientation similar to that
of tetrahedral shape
Similar to arrangement in
tetrahedral
Similar to arrangement in
tetrahedral
of ammonia NH3 andwater H2O
Shape and angleShape trigonal pyramidal
Angle 1070
Shape bent
Angle 10450
Number of bond pair amp
lone pair electrons
Bond pair electron 3
Lone pair electron 1
Bond pair electron 2
Lone pair electron 2
bull The angle of bond pair - bond pair electrons in ammonia and0 0
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water are 1070 and 10450 respectively which is lesser than in
methane molecule (10950
) This can be explained by the factof the presence of lone-pair electrons in both ammonia and
water Since the lone pair - lone pair electron repulsion is
stronger than lone pair - bond pair electron repulsion than
bond pair - bond pair electron repulsion it is expected that therepulsion between lone pair - bond pair electrons in ammonia
is stronger than bond pair - bond pair electrons repulsion
hence caused the angle to squeeze to a smaller angle As
for water since there is a presence of lone pair - lone pair
electrons repulsion it results the bond pair - bond pair
electron repulsion to be much smaller hence caused the
angle to squeeze to a smaller angle Though valence bondtheory does not actually explain the hybridisation especially
when it relates to the repulsion occur involving lone pair
electron another theories shall be applied to study such
effect All these shall be discussed further in VSEPR theory
b) Effect of electronegativity and bonding angle between bond pair
- bond pair electrons ~
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- bond pair electrons
i As discussed earlier in the bonding of water (H2
O) the
molecular shape and angle is described in the diagram below
Sul hur which is also an element from Grou 16 formed h dro en
sulphide H2S when sulphur react with hydrogen However unlike
water the bonding angles is much smaller compare to water This is
due to the difference of electronegativity and also the bond length
between O and S in molecule Since O is more electronegative
compare to S the O-H bond is pulled closer toward O Furthermorethe bond length of O-H is shorter compare to S-H As a result the
bonding pair-bonding pair electrons repulse greater with each other in
H-O-H and caused the angle become greater
ii Another example is between NH3 and PH3 The orbital
f
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diagram for both NH3 and PH3 are described below Both
nitrogen N and phosphorous P are from the same groupwhich is Group 15
Similar to the above case P in phosphine is less
electronegative than N in ammonia and bond length of N-H
is shorter than P-H As a result H is pull closer to N and
repulsion between bonding of H-N-H is greater compare toH-P-H and caused the angle between H-N-H is greater
compare to H-P-H
324 Valence Shell Electron Pair Repulsion (VSEPR) Theory
1 M l l t i th th di i l t f t
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1 Molecular geometry is the three-dimensional arrangement of atoms
in a molecule A moleculersquos geometry affects its physical andchemical properties such as melting point boiling point density
and the types of reactions it undergoes The basic concept of
VSEPR theory is based on the three general rules below
a) As far as electron-pair repulsion is concerned double bonds andtriple bonds can be treated like single bonds This approximation is
good for qualitative purposes However you should realize that in
ldquo rdquo
because there are two or three bonds between two atoms theelectron density occupies more space
b) If a molecule has two or more resonance structures we can apply
the VSEPR model to any one of them Formal charges are usually
not shown
c) The order of repulsion strength of lone pair and bond pair are
lone-pair amp lone-pair electrons repulsion are the strongest
followed by lone-pair amp bond-pair electrons repulsion while
bond-pair amp bond-pair electrons repulsion is the weakest
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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diagram
a) Ground state
Both nitrogen whichact as the center atom
has the valence
electron of 2s22p3
b) Excited state One of the electron
from 2s is promote to 2p
orbitals - e ual ener
level
c) Hybridised state
One orbital of 2s and
one orbitals of 2p
combined (hybrid) and
rearrange themselves
to the shape and
orientation of a linear
shapeMolecular shape Linear
Angle between bond pair 1800
b) Whereas for hydrogen cyanide HCN both carbon atom and
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) y g y
nitrogen atom undergoes sp hybridisation in order to form a
linear structureExplanation
Energy level diagram for
carbon atom C
Energy level diagram for
nitrogen atom N
a) Ground state
Valence electronC 2s22p2
N 2s22p3
b) Excited state
One of the electron from
2s is promote to 2p
orbitals - equal energy
level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined
(hybrid) and form sp
hybridisation
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N
σσσσ
π
σσσσ
π
6 However there are a few limitation on valence bond theory such
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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as when explaining the effect angle of bond-pair and bond-pair
electrons when there isare presence of lone pair electron in thecenter atom and also the difference of electronegativity
a) Bonding in ammonia NH3 and water H2O ~ Both NH3 and
H2O undergoes an arrangement similar to sp3 hybridisation
similar to that of C in methane Table below shows thehybridisation occur for nitrogen in ammonia and oxygen in water
Energy level diagram Energy level diagram
for nitrogen atom N for oxygen atom O
a) Ground state
Valence electron
N 2s22p3
O 2s2
2p4
b) Excited state
One of the electron
from 2s is promote to
2p orbitals
c) Hybridised state
One orbital of 2s and
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One orbital of 2s and
three orbitals from 2p
combined (hybrid) andrearrange themselves to
the shape and
orientation similar to that
of tetrahedral shape
Similar to arrangement in
tetrahedral
Similar to arrangement in
tetrahedral
of ammonia NH3 andwater H2O
Shape and angleShape trigonal pyramidal
Angle 1070
Shape bent
Angle 10450
Number of bond pair amp
lone pair electrons
Bond pair electron 3
Lone pair electron 1
Bond pair electron 2
Lone pair electron 2
bull The angle of bond pair - bond pair electrons in ammonia and0 0
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water are 1070 and 10450 respectively which is lesser than in
methane molecule (10950
) This can be explained by the factof the presence of lone-pair electrons in both ammonia and
water Since the lone pair - lone pair electron repulsion is
stronger than lone pair - bond pair electron repulsion than
bond pair - bond pair electron repulsion it is expected that therepulsion between lone pair - bond pair electrons in ammonia
is stronger than bond pair - bond pair electrons repulsion
hence caused the angle to squeeze to a smaller angle As
for water since there is a presence of lone pair - lone pair
electrons repulsion it results the bond pair - bond pair
electron repulsion to be much smaller hence caused the
angle to squeeze to a smaller angle Though valence bondtheory does not actually explain the hybridisation especially
when it relates to the repulsion occur involving lone pair
electron another theories shall be applied to study such
effect All these shall be discussed further in VSEPR theory
b) Effect of electronegativity and bonding angle between bond pair
- bond pair electrons ~
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- bond pair electrons
i As discussed earlier in the bonding of water (H2
O) the
molecular shape and angle is described in the diagram below
Sul hur which is also an element from Grou 16 formed h dro en
sulphide H2S when sulphur react with hydrogen However unlike
water the bonding angles is much smaller compare to water This is
due to the difference of electronegativity and also the bond length
between O and S in molecule Since O is more electronegative
compare to S the O-H bond is pulled closer toward O Furthermorethe bond length of O-H is shorter compare to S-H As a result the
bonding pair-bonding pair electrons repulse greater with each other in
H-O-H and caused the angle become greater
ii Another example is between NH3 and PH3 The orbital
f
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diagram for both NH3 and PH3 are described below Both
nitrogen N and phosphorous P are from the same groupwhich is Group 15
Similar to the above case P in phosphine is less
electronegative than N in ammonia and bond length of N-H
is shorter than P-H As a result H is pull closer to N and
repulsion between bonding of H-N-H is greater compare toH-P-H and caused the angle between H-N-H is greater
compare to H-P-H
324 Valence Shell Electron Pair Repulsion (VSEPR) Theory
1 M l l t i th th di i l t f t
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1 Molecular geometry is the three-dimensional arrangement of atoms
in a molecule A moleculersquos geometry affects its physical andchemical properties such as melting point boiling point density
and the types of reactions it undergoes The basic concept of
VSEPR theory is based on the three general rules below
a) As far as electron-pair repulsion is concerned double bonds andtriple bonds can be treated like single bonds This approximation is
good for qualitative purposes However you should realize that in
ldquo rdquo
because there are two or three bonds between two atoms theelectron density occupies more space
b) If a molecule has two or more resonance structures we can apply
the VSEPR model to any one of them Formal charges are usually
not shown
c) The order of repulsion strength of lone pair and bond pair are
lone-pair amp lone-pair electrons repulsion are the strongest
followed by lone-pair amp bond-pair electrons repulsion while
bond-pair amp bond-pair electrons repulsion is the weakest
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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) y g y
nitrogen atom undergoes sp hybridisation in order to form a
linear structureExplanation
Energy level diagram for
carbon atom C
Energy level diagram for
nitrogen atom N
a) Ground state
Valence electronC 2s22p2
N 2s22p3
b) Excited state
One of the electron from
2s is promote to 2p
orbitals - equal energy
level
c) Hybridised state
One orbital of 2s and one
orbitals of 2p combined
(hybrid) and form sp
hybridisation
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N
σσσσ
π
σσσσ
π
6 However there are a few limitation on valence bond theory such
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as when explaining the effect angle of bond-pair and bond-pair
electrons when there isare presence of lone pair electron in thecenter atom and also the difference of electronegativity
a) Bonding in ammonia NH3 and water H2O ~ Both NH3 and
H2O undergoes an arrangement similar to sp3 hybridisation
similar to that of C in methane Table below shows thehybridisation occur for nitrogen in ammonia and oxygen in water
Energy level diagram Energy level diagram
for nitrogen atom N for oxygen atom O
a) Ground state
Valence electron
N 2s22p3
O 2s2
2p4
b) Excited state
One of the electron
from 2s is promote to
2p orbitals
c) Hybridised state
One orbital of 2s and
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One orbital of 2s and
three orbitals from 2p
combined (hybrid) andrearrange themselves to
the shape and
orientation similar to that
of tetrahedral shape
Similar to arrangement in
tetrahedral
Similar to arrangement in
tetrahedral
of ammonia NH3 andwater H2O
Shape and angleShape trigonal pyramidal
Angle 1070
Shape bent
Angle 10450
Number of bond pair amp
lone pair electrons
Bond pair electron 3
Lone pair electron 1
Bond pair electron 2
Lone pair electron 2
bull The angle of bond pair - bond pair electrons in ammonia and0 0
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water are 1070 and 10450 respectively which is lesser than in
methane molecule (10950
) This can be explained by the factof the presence of lone-pair electrons in both ammonia and
water Since the lone pair - lone pair electron repulsion is
stronger than lone pair - bond pair electron repulsion than
bond pair - bond pair electron repulsion it is expected that therepulsion between lone pair - bond pair electrons in ammonia
is stronger than bond pair - bond pair electrons repulsion
hence caused the angle to squeeze to a smaller angle As
for water since there is a presence of lone pair - lone pair
electrons repulsion it results the bond pair - bond pair
electron repulsion to be much smaller hence caused the
angle to squeeze to a smaller angle Though valence bondtheory does not actually explain the hybridisation especially
when it relates to the repulsion occur involving lone pair
electron another theories shall be applied to study such
effect All these shall be discussed further in VSEPR theory
b) Effect of electronegativity and bonding angle between bond pair
- bond pair electrons ~
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- bond pair electrons
i As discussed earlier in the bonding of water (H2
O) the
molecular shape and angle is described in the diagram below
Sul hur which is also an element from Grou 16 formed h dro en
sulphide H2S when sulphur react with hydrogen However unlike
water the bonding angles is much smaller compare to water This is
due to the difference of electronegativity and also the bond length
between O and S in molecule Since O is more electronegative
compare to S the O-H bond is pulled closer toward O Furthermorethe bond length of O-H is shorter compare to S-H As a result the
bonding pair-bonding pair electrons repulse greater with each other in
H-O-H and caused the angle become greater
ii Another example is between NH3 and PH3 The orbital
f
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diagram for both NH3 and PH3 are described below Both
nitrogen N and phosphorous P are from the same groupwhich is Group 15
Similar to the above case P in phosphine is less
electronegative than N in ammonia and bond length of N-H
is shorter than P-H As a result H is pull closer to N and
repulsion between bonding of H-N-H is greater compare toH-P-H and caused the angle between H-N-H is greater
compare to H-P-H
324 Valence Shell Electron Pair Repulsion (VSEPR) Theory
1 M l l t i th th di i l t f t
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1 Molecular geometry is the three-dimensional arrangement of atoms
in a molecule A moleculersquos geometry affects its physical andchemical properties such as melting point boiling point density
and the types of reactions it undergoes The basic concept of
VSEPR theory is based on the three general rules below
a) As far as electron-pair repulsion is concerned double bonds andtriple bonds can be treated like single bonds This approximation is
good for qualitative purposes However you should realize that in
ldquo rdquo
because there are two or three bonds between two atoms theelectron density occupies more space
b) If a molecule has two or more resonance structures we can apply
the VSEPR model to any one of them Formal charges are usually
not shown
c) The order of repulsion strength of lone pair and bond pair are
lone-pair amp lone-pair electrons repulsion are the strongest
followed by lone-pair amp bond-pair electrons repulsion while
bond-pair amp bond-pair electrons repulsion is the weakest
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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N
σσσσ
π
σσσσ
π
6 However there are a few limitation on valence bond theory such
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as when explaining the effect angle of bond-pair and bond-pair
electrons when there isare presence of lone pair electron in thecenter atom and also the difference of electronegativity
a) Bonding in ammonia NH3 and water H2O ~ Both NH3 and
H2O undergoes an arrangement similar to sp3 hybridisation
similar to that of C in methane Table below shows thehybridisation occur for nitrogen in ammonia and oxygen in water
Energy level diagram Energy level diagram
for nitrogen atom N for oxygen atom O
a) Ground state
Valence electron
N 2s22p3
O 2s2
2p4
b) Excited state
One of the electron
from 2s is promote to
2p orbitals
c) Hybridised state
One orbital of 2s and
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One orbital of 2s and
three orbitals from 2p
combined (hybrid) andrearrange themselves to
the shape and
orientation similar to that
of tetrahedral shape
Similar to arrangement in
tetrahedral
Similar to arrangement in
tetrahedral
of ammonia NH3 andwater H2O
Shape and angleShape trigonal pyramidal
Angle 1070
Shape bent
Angle 10450
Number of bond pair amp
lone pair electrons
Bond pair electron 3
Lone pair electron 1
Bond pair electron 2
Lone pair electron 2
bull The angle of bond pair - bond pair electrons in ammonia and0 0
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water are 1070 and 10450 respectively which is lesser than in
methane molecule (10950
) This can be explained by the factof the presence of lone-pair electrons in both ammonia and
water Since the lone pair - lone pair electron repulsion is
stronger than lone pair - bond pair electron repulsion than
bond pair - bond pair electron repulsion it is expected that therepulsion between lone pair - bond pair electrons in ammonia
is stronger than bond pair - bond pair electrons repulsion
hence caused the angle to squeeze to a smaller angle As
for water since there is a presence of lone pair - lone pair
electrons repulsion it results the bond pair - bond pair
electron repulsion to be much smaller hence caused the
angle to squeeze to a smaller angle Though valence bondtheory does not actually explain the hybridisation especially
when it relates to the repulsion occur involving lone pair
electron another theories shall be applied to study such
effect All these shall be discussed further in VSEPR theory
b) Effect of electronegativity and bonding angle between bond pair
- bond pair electrons ~
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- bond pair electrons
i As discussed earlier in the bonding of water (H2
O) the
molecular shape and angle is described in the diagram below
Sul hur which is also an element from Grou 16 formed h dro en
sulphide H2S when sulphur react with hydrogen However unlike
water the bonding angles is much smaller compare to water This is
due to the difference of electronegativity and also the bond length
between O and S in molecule Since O is more electronegative
compare to S the O-H bond is pulled closer toward O Furthermorethe bond length of O-H is shorter compare to S-H As a result the
bonding pair-bonding pair electrons repulse greater with each other in
H-O-H and caused the angle become greater
ii Another example is between NH3 and PH3 The orbital
f
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diagram for both NH3 and PH3 are described below Both
nitrogen N and phosphorous P are from the same groupwhich is Group 15
Similar to the above case P in phosphine is less
electronegative than N in ammonia and bond length of N-H
is shorter than P-H As a result H is pull closer to N and
repulsion between bonding of H-N-H is greater compare toH-P-H and caused the angle between H-N-H is greater
compare to H-P-H
324 Valence Shell Electron Pair Repulsion (VSEPR) Theory
1 M l l t i th th di i l t f t
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1 Molecular geometry is the three-dimensional arrangement of atoms
in a molecule A moleculersquos geometry affects its physical andchemical properties such as melting point boiling point density
and the types of reactions it undergoes The basic concept of
VSEPR theory is based on the three general rules below
a) As far as electron-pair repulsion is concerned double bonds andtriple bonds can be treated like single bonds This approximation is
good for qualitative purposes However you should realize that in
ldquo rdquo
because there are two or three bonds between two atoms theelectron density occupies more space
b) If a molecule has two or more resonance structures we can apply
the VSEPR model to any one of them Formal charges are usually
not shown
c) The order of repulsion strength of lone pair and bond pair are
lone-pair amp lone-pair electrons repulsion are the strongest
followed by lone-pair amp bond-pair electrons repulsion while
bond-pair amp bond-pair electrons repulsion is the weakest
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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as when explaining the effect angle of bond-pair and bond-pair
electrons when there isare presence of lone pair electron in thecenter atom and also the difference of electronegativity
a) Bonding in ammonia NH3 and water H2O ~ Both NH3 and
H2O undergoes an arrangement similar to sp3 hybridisation
similar to that of C in methane Table below shows thehybridisation occur for nitrogen in ammonia and oxygen in water
Energy level diagram Energy level diagram
for nitrogen atom N for oxygen atom O
a) Ground state
Valence electron
N 2s22p3
O 2s2
2p4
b) Excited state
One of the electron
from 2s is promote to
2p orbitals
c) Hybridised state
One orbital of 2s and
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One orbital of 2s and
three orbitals from 2p
combined (hybrid) andrearrange themselves to
the shape and
orientation similar to that
of tetrahedral shape
Similar to arrangement in
tetrahedral
Similar to arrangement in
tetrahedral
of ammonia NH3 andwater H2O
Shape and angleShape trigonal pyramidal
Angle 1070
Shape bent
Angle 10450
Number of bond pair amp
lone pair electrons
Bond pair electron 3
Lone pair electron 1
Bond pair electron 2
Lone pair electron 2
bull The angle of bond pair - bond pair electrons in ammonia and0 0
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water are 1070 and 10450 respectively which is lesser than in
methane molecule (10950
) This can be explained by the factof the presence of lone-pair electrons in both ammonia and
water Since the lone pair - lone pair electron repulsion is
stronger than lone pair - bond pair electron repulsion than
bond pair - bond pair electron repulsion it is expected that therepulsion between lone pair - bond pair electrons in ammonia
is stronger than bond pair - bond pair electrons repulsion
hence caused the angle to squeeze to a smaller angle As
for water since there is a presence of lone pair - lone pair
electrons repulsion it results the bond pair - bond pair
electron repulsion to be much smaller hence caused the
angle to squeeze to a smaller angle Though valence bondtheory does not actually explain the hybridisation especially
when it relates to the repulsion occur involving lone pair
electron another theories shall be applied to study such
effect All these shall be discussed further in VSEPR theory
b) Effect of electronegativity and bonding angle between bond pair
- bond pair electrons ~
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- bond pair electrons
i As discussed earlier in the bonding of water (H2
O) the
molecular shape and angle is described in the diagram below
Sul hur which is also an element from Grou 16 formed h dro en
sulphide H2S when sulphur react with hydrogen However unlike
water the bonding angles is much smaller compare to water This is
due to the difference of electronegativity and also the bond length
between O and S in molecule Since O is more electronegative
compare to S the O-H bond is pulled closer toward O Furthermorethe bond length of O-H is shorter compare to S-H As a result the
bonding pair-bonding pair electrons repulse greater with each other in
H-O-H and caused the angle become greater
ii Another example is between NH3 and PH3 The orbital
f
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diagram for both NH3 and PH3 are described below Both
nitrogen N and phosphorous P are from the same groupwhich is Group 15
Similar to the above case P in phosphine is less
electronegative than N in ammonia and bond length of N-H
is shorter than P-H As a result H is pull closer to N and
repulsion between bonding of H-N-H is greater compare toH-P-H and caused the angle between H-N-H is greater
compare to H-P-H
324 Valence Shell Electron Pair Repulsion (VSEPR) Theory
1 M l l t i th th di i l t f t
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1 Molecular geometry is the three-dimensional arrangement of atoms
in a molecule A moleculersquos geometry affects its physical andchemical properties such as melting point boiling point density
and the types of reactions it undergoes The basic concept of
VSEPR theory is based on the three general rules below
a) As far as electron-pair repulsion is concerned double bonds andtriple bonds can be treated like single bonds This approximation is
good for qualitative purposes However you should realize that in
ldquo rdquo
because there are two or three bonds between two atoms theelectron density occupies more space
b) If a molecule has two or more resonance structures we can apply
the VSEPR model to any one of them Formal charges are usually
not shown
c) The order of repulsion strength of lone pair and bond pair are
lone-pair amp lone-pair electrons repulsion are the strongest
followed by lone-pair amp bond-pair electrons repulsion while
bond-pair amp bond-pair electrons repulsion is the weakest
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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One orbital of 2s and
three orbitals from 2p
combined (hybrid) andrearrange themselves to
the shape and
orientation similar to that
of tetrahedral shape
Similar to arrangement in
tetrahedral
Similar to arrangement in
tetrahedral
of ammonia NH3 andwater H2O
Shape and angleShape trigonal pyramidal
Angle 1070
Shape bent
Angle 10450
Number of bond pair amp
lone pair electrons
Bond pair electron 3
Lone pair electron 1
Bond pair electron 2
Lone pair electron 2
bull The angle of bond pair - bond pair electrons in ammonia and0 0
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water are 1070 and 10450 respectively which is lesser than in
methane molecule (10950
) This can be explained by the factof the presence of lone-pair electrons in both ammonia and
water Since the lone pair - lone pair electron repulsion is
stronger than lone pair - bond pair electron repulsion than
bond pair - bond pair electron repulsion it is expected that therepulsion between lone pair - bond pair electrons in ammonia
is stronger than bond pair - bond pair electrons repulsion
hence caused the angle to squeeze to a smaller angle As
for water since there is a presence of lone pair - lone pair
electrons repulsion it results the bond pair - bond pair
electron repulsion to be much smaller hence caused the
angle to squeeze to a smaller angle Though valence bondtheory does not actually explain the hybridisation especially
when it relates to the repulsion occur involving lone pair
electron another theories shall be applied to study such
effect All these shall be discussed further in VSEPR theory
b) Effect of electronegativity and bonding angle between bond pair
- bond pair electrons ~
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- bond pair electrons
i As discussed earlier in the bonding of water (H2
O) the
molecular shape and angle is described in the diagram below
Sul hur which is also an element from Grou 16 formed h dro en
sulphide H2S when sulphur react with hydrogen However unlike
water the bonding angles is much smaller compare to water This is
due to the difference of electronegativity and also the bond length
between O and S in molecule Since O is more electronegative
compare to S the O-H bond is pulled closer toward O Furthermorethe bond length of O-H is shorter compare to S-H As a result the
bonding pair-bonding pair electrons repulse greater with each other in
H-O-H and caused the angle become greater
ii Another example is between NH3 and PH3 The orbital
f
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diagram for both NH3 and PH3 are described below Both
nitrogen N and phosphorous P are from the same groupwhich is Group 15
Similar to the above case P in phosphine is less
electronegative than N in ammonia and bond length of N-H
is shorter than P-H As a result H is pull closer to N and
repulsion between bonding of H-N-H is greater compare toH-P-H and caused the angle between H-N-H is greater
compare to H-P-H
324 Valence Shell Electron Pair Repulsion (VSEPR) Theory
1 M l l t i th th di i l t f t
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1 Molecular geometry is the three-dimensional arrangement of atoms
in a molecule A moleculersquos geometry affects its physical andchemical properties such as melting point boiling point density
and the types of reactions it undergoes The basic concept of
VSEPR theory is based on the three general rules below
a) As far as electron-pair repulsion is concerned double bonds andtriple bonds can be treated like single bonds This approximation is
good for qualitative purposes However you should realize that in
ldquo rdquo
because there are two or three bonds between two atoms theelectron density occupies more space
b) If a molecule has two or more resonance structures we can apply
the VSEPR model to any one of them Formal charges are usually
not shown
c) The order of repulsion strength of lone pair and bond pair are
lone-pair amp lone-pair electrons repulsion are the strongest
followed by lone-pair amp bond-pair electrons repulsion while
bond-pair amp bond-pair electrons repulsion is the weakest
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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water are 1070 and 10450 respectively which is lesser than in
methane molecule (10950
) This can be explained by the factof the presence of lone-pair electrons in both ammonia and
water Since the lone pair - lone pair electron repulsion is
stronger than lone pair - bond pair electron repulsion than
bond pair - bond pair electron repulsion it is expected that therepulsion between lone pair - bond pair electrons in ammonia
is stronger than bond pair - bond pair electrons repulsion
hence caused the angle to squeeze to a smaller angle As
for water since there is a presence of lone pair - lone pair
electrons repulsion it results the bond pair - bond pair
electron repulsion to be much smaller hence caused the
angle to squeeze to a smaller angle Though valence bondtheory does not actually explain the hybridisation especially
when it relates to the repulsion occur involving lone pair
electron another theories shall be applied to study such
effect All these shall be discussed further in VSEPR theory
b) Effect of electronegativity and bonding angle between bond pair
- bond pair electrons ~
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- bond pair electrons
i As discussed earlier in the bonding of water (H2
O) the
molecular shape and angle is described in the diagram below
Sul hur which is also an element from Grou 16 formed h dro en
sulphide H2S when sulphur react with hydrogen However unlike
water the bonding angles is much smaller compare to water This is
due to the difference of electronegativity and also the bond length
between O and S in molecule Since O is more electronegative
compare to S the O-H bond is pulled closer toward O Furthermorethe bond length of O-H is shorter compare to S-H As a result the
bonding pair-bonding pair electrons repulse greater with each other in
H-O-H and caused the angle become greater
ii Another example is between NH3 and PH3 The orbital
f
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diagram for both NH3 and PH3 are described below Both
nitrogen N and phosphorous P are from the same groupwhich is Group 15
Similar to the above case P in phosphine is less
electronegative than N in ammonia and bond length of N-H
is shorter than P-H As a result H is pull closer to N and
repulsion between bonding of H-N-H is greater compare toH-P-H and caused the angle between H-N-H is greater
compare to H-P-H
324 Valence Shell Electron Pair Repulsion (VSEPR) Theory
1 M l l t i th th di i l t f t
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1 Molecular geometry is the three-dimensional arrangement of atoms
in a molecule A moleculersquos geometry affects its physical andchemical properties such as melting point boiling point density
and the types of reactions it undergoes The basic concept of
VSEPR theory is based on the three general rules below
a) As far as electron-pair repulsion is concerned double bonds andtriple bonds can be treated like single bonds This approximation is
good for qualitative purposes However you should realize that in
ldquo rdquo
because there are two or three bonds between two atoms theelectron density occupies more space
b) If a molecule has two or more resonance structures we can apply
the VSEPR model to any one of them Formal charges are usually
not shown
c) The order of repulsion strength of lone pair and bond pair are
lone-pair amp lone-pair electrons repulsion are the strongest
followed by lone-pair amp bond-pair electrons repulsion while
bond-pair amp bond-pair electrons repulsion is the weakest
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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- bond pair electrons
i As discussed earlier in the bonding of water (H2
O) the
molecular shape and angle is described in the diagram below
Sul hur which is also an element from Grou 16 formed h dro en
sulphide H2S when sulphur react with hydrogen However unlike
water the bonding angles is much smaller compare to water This is
due to the difference of electronegativity and also the bond length
between O and S in molecule Since O is more electronegative
compare to S the O-H bond is pulled closer toward O Furthermorethe bond length of O-H is shorter compare to S-H As a result the
bonding pair-bonding pair electrons repulse greater with each other in
H-O-H and caused the angle become greater
ii Another example is between NH3 and PH3 The orbital
f
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diagram for both NH3 and PH3 are described below Both
nitrogen N and phosphorous P are from the same groupwhich is Group 15
Similar to the above case P in phosphine is less
electronegative than N in ammonia and bond length of N-H
is shorter than P-H As a result H is pull closer to N and
repulsion between bonding of H-N-H is greater compare toH-P-H and caused the angle between H-N-H is greater
compare to H-P-H
324 Valence Shell Electron Pair Repulsion (VSEPR) Theory
1 M l l t i th th di i l t f t
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1 Molecular geometry is the three-dimensional arrangement of atoms
in a molecule A moleculersquos geometry affects its physical andchemical properties such as melting point boiling point density
and the types of reactions it undergoes The basic concept of
VSEPR theory is based on the three general rules below
a) As far as electron-pair repulsion is concerned double bonds andtriple bonds can be treated like single bonds This approximation is
good for qualitative purposes However you should realize that in
ldquo rdquo
because there are two or three bonds between two atoms theelectron density occupies more space
b) If a molecule has two or more resonance structures we can apply
the VSEPR model to any one of them Formal charges are usually
not shown
c) The order of repulsion strength of lone pair and bond pair are
lone-pair amp lone-pair electrons repulsion are the strongest
followed by lone-pair amp bond-pair electrons repulsion while
bond-pair amp bond-pair electrons repulsion is the weakest
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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diagram for both NH3 and PH3 are described below Both
nitrogen N and phosphorous P are from the same groupwhich is Group 15
Similar to the above case P in phosphine is less
electronegative than N in ammonia and bond length of N-H
is shorter than P-H As a result H is pull closer to N and
repulsion between bonding of H-N-H is greater compare toH-P-H and caused the angle between H-N-H is greater
compare to H-P-H
324 Valence Shell Electron Pair Repulsion (VSEPR) Theory
1 M l l t i th th di i l t f t
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1 Molecular geometry is the three-dimensional arrangement of atoms
in a molecule A moleculersquos geometry affects its physical andchemical properties such as melting point boiling point density
and the types of reactions it undergoes The basic concept of
VSEPR theory is based on the three general rules below
a) As far as electron-pair repulsion is concerned double bonds andtriple bonds can be treated like single bonds This approximation is
good for qualitative purposes However you should realize that in
ldquo rdquo
because there are two or three bonds between two atoms theelectron density occupies more space
b) If a molecule has two or more resonance structures we can apply
the VSEPR model to any one of them Formal charges are usually
not shown
c) The order of repulsion strength of lone pair and bond pair are
lone-pair amp lone-pair electrons repulsion are the strongest
followed by lone-pair amp bond-pair electrons repulsion while
bond-pair amp bond-pair electrons repulsion is the weakest
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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1 Molecular geometry is the three-dimensional arrangement of atoms
in a molecule A moleculersquos geometry affects its physical andchemical properties such as melting point boiling point density
and the types of reactions it undergoes The basic concept of
VSEPR theory is based on the three general rules below
a) As far as electron-pair repulsion is concerned double bonds andtriple bonds can be treated like single bonds This approximation is
good for qualitative purposes However you should realize that in
ldquo rdquo
because there are two or three bonds between two atoms theelectron density occupies more space
b) If a molecule has two or more resonance structures we can apply
the VSEPR model to any one of them Formal charges are usually
not shown
c) The order of repulsion strength of lone pair and bond pair are
lone-pair amp lone-pair electrons repulsion are the strongest
followed by lone-pair amp bond-pair electrons repulsion while
bond-pair amp bond-pair electrons repulsion is the weakest
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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No of
surroun
No of
loneMolecular Diagram of the molecular Example of
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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Classsurroun
d
atoms
pair
electron
Molecular
geometry
Diagram of the molecular
shape
Example of
molecules
AB2 2 0 Linear CO2
BeCl2
AB3 3 0Trigonal
lanar
AlCl3BF3
NO3-
AB2E 2 1Shape
Bent
SO2
O3
NO2-
AB4 4 0Shape
Tetrahedral
CH4
SiCl4SO
4
2-
Shape NH3
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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AB3E 3 1 Trigonal
pyramidal
PCl3
SO32-
AB2E2 2 2 Shape Bent
H2O
SCl2H2O2
AB5 5 0
Shape
Trigonal
bipyramidal
PCl5SbF5
AB4E 4 1Shape
See-Saw
SCl4PF4
-
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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AB3E2 3 2
Arrangeme
nt
Trigonal
bipyramidalShape
T-shape
ICl3
BrF3
AB2E3 2 3
Arrangement
Trigonal
bipyramidal
Shape
linear
I3-
BrCl2-
A
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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AB6 6 0
Arrangeme
ntamp Shape
Octahedral
SF6
AB5E 5 1
Arrangement
Octahedral
Sha e
SbCl52-
IF
Square
pyramidal
AB4E2 4 2
Arrangeme
nt
OctahedralShape
Square
planar
XeF4
BrF4-
a) phosphorous trichloride PCl3S1 Total valence electrons
P = 5 e- 3 Cl = 3 x 7e-
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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P = 5 e- 3 Cl = 3 x 7e
Total electrons = 26
S2 Electrons used bond = 3 x 2e-
Electrons left = 26 - 6 = 20 e-
S3 e- at surround atom = 3 x 6e-Electrons left = 20 - 18 = 2 e-
S4a ) The 2e- remain is placed atthe center atom P
Since molecule contain 3
surrounding atom and 1 lone pairelectrons hence
Arrangement tetrahedral
Shape trigonal pyramidal
b) Carbonate ion CO32-
S1 Total valence electrons
C = 4 e 3 O = 3 x 6e- + 2e accept
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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C = 4 e- 3 O = 3 x 6e + 2e- accept Total electrons = 24
S2 Electrons used bond = 3 x 2e-
Electrons left = 24 - 6 = 18 e-
- -
Electrons left = 18 - 18 = 0 e-
S4b) Since the center atom C not
yet achieved octet a double bondis form using any e- from O
Since molecular ion contain 3
surrounding atom and 0 lone pairelectrons hence
Arrangement and
Shape trigonal planar
c) iodine tetrachloride ion ICl4-
S1 Total valence electrons
I = 7 e- 4 Cl = 4 x 7e- + 1e- accept
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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I = 7 e 4 Cl = 4 x 7e + 1e accept Total electrons = 36
S2 Electrons used bond = 4 x 2e-
Electrons left = 36 - 8 = 28 e-
- = -
Electrons left = 28 - 24 = 4 e-
S4a ) The 2e- remain is placed at
the center atom P
Since molecular ion contain 4
surrounding atom and 2 lone pairelectrons hence
Arrangement octahedralShape square planar
a) Iodide ion I3- b) Antimony pentachloride ion
[SbCl ]2
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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[SbCl5]2-
Arrangement trigonal
bipyramidalShape linear
Arrangement octahedralShape square pyramidal
3 4
Arrangement trigonalbipyramidal
Shape T-shape
Arrangement trigonalbipyramidal
Shape see-saw
325 Electronegativity and Polarity of Molecules
1 From all the chemical bonding discussed so far ionic and
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7499
polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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o a e c e ca bo d g d scussed so a o c a d
covalent bonding models portray compounds as beingformed by either complete electron transfer or complete
electron sharing However in most real compounds the
type of bonding lies somewhere between these extremes
Thus the great majority of bonds are more accuratelythought of as ldquopolar covalentrdquo that is partially ionic and
partially covalent
Pure ionic compound Pure covalent compound Polar covalent compound
2 One of the most important concepts in chemical bonding is
electronegativity (EN) the relative ability of a bonded atom to
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7499
polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7699
whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7799
present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8199
b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8299
c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8399
usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8699
34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8899
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9199
the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9399
example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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attract the shared electrons Electronegativity is a relative concept
meaning that an elementrsquos electronegativity can be measured only
in relation to the electronegativity of other elements Linus Pauling
devised a method for calculating relative electronegativities of most
elements
Molecule Fluorine F2 Hydrogen fluoride HF
δ+ δ
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 6899
with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7399
more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7499
polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7599
density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7699
whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7799
present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7899
983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7999
charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8099
Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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Lewis
structure
δ+ δminus
Polarity Non-polar molecule Polar molecule
∆EN 40 - 40 = 0 40 - 21 = 19
Since there are no
different between the EN
Since F is more electronegative than H
therefore the bondin air electrons
Explanation
therefore bonding pair
electrons was not pulled to
either atom hence remain
in the middle between 2 F
atom
were pulled closer to F atom This will
caused F to have greater electron
density compare to H Therefore F has
partial negative charge (δminus) while H
carries partial positive charge (δ+)
Dipole
moment
magnitude
and
vector
Since there is no
difference between EN
the dipole moment is 0
and no resultant dipole
moment nor vector
Since F is more electronegative than H
There is presence of dipole moment in
HF and the vector of resultant dipole
moment is pointed to the direction of F
(symbolised by I )
a) Comparisons above are basically the difference between an
element with compound where diatomic molecules
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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containing atoms of different elements (for example HClCO and NO) have dipole moments and are called polar
molecules while diatomic molecules containing atoms of
the same element (for example H2 O2 and F2) are
examples of non-polar molecules because they do nothave dipole moments However not necessarily a covalent
bond compound is guaranteed a polar molecule
b) For a molecule made up of three or more atoms both the
polarity of the bonds and the molecular geometry determine
whether there is a dipole moment Even if polar bonds are
present the molecule will not necessarily have a dipole
moment For example comparison between sulphur dioxideand sulphur trioxide
Molecules Sulphur trioxide SO3 Sulphur dioxide SO2
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
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example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
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water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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Lewis
structure
and shape
Polarity Non-polar molecule Polar molecule
EN S = 25 O = 35 S = 25 O = 35
the dipole moment of the entire molecule
is made u of three bond moments that isthe dipole moment of the entire molecule
Bond
moment
dipole
moment
magnitude
and vector
individual dipole moments in the polar
SminusO bonds The bond moment is a vector
quantity which means that it has both
magnitude and direction The measured
dipole moment is equal to the vector sum
of the bond moments The three bondmoments in SO3 are equal in magnitude
Because they point in opposite directions
in a planar SO3 molecule the sum of
resultant dipole moment would be zero
lone pair electron Even though two bond
moments in SO2 are equal in magnitude
however the presence of the lone pair
electrons which caused the repulsion of
bond pair electrons to be lesser Because
they point in downward directions in a bent
SO2 molecule the overall vector points
downward and the sum of resultant
dipole moment would not be zero
hence a polar molecule
From the example of SO2 and SO3 used we can tell that if a
polyatomic molecule is a symmetrical molecule (molecule
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7699
whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7799
present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7899
983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8099
Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8199
b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8299
c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8399
usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8499
3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8599
substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8699
34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8799
are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8899
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8999
3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9099
Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9199
the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9399
example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9499
dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9799
Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9899
water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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with no lone pair electrons in it) it may be a non-polarmolecule However if a polyatomic molecule is an
asymmetrical molecule (molecule with lone pair electrons in
it) it may be a polar molecules
c) Even though a polyatomic molecule may be symmetrical if
polar molecule as the bonding moments are different andcaused the magnitude of dipole moment of the molecule is not
equal to zero However if the surrounding atoms are the
same bonding moments are equal in magnitude and the
resultant vector cancel-off each other causing the dipolemoment is equal to zero hence form a non-polar molecule
For example
MoleculeMethane CH4 Chloromethane
( hl f ) CH Cl
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7499
polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7599
density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7699
whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7799
present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7899
983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7999
charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8099
Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8199
b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8299
c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8399
usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8699
34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8799
are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8899
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9099
Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9199
the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9399
example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9799
Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9899
water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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(chloroform) CH3Cl
Lewis
structure
-
Explanation
As methane is a symmetrical
molecule and the surrounding atoms
are the same the vector of bond
moment cancel off each other hencecaused the dipole moment is equal to
zero
Since a foreign element Cl is in thesymmetrical molecule and Cl is more
EN than the rest of the atoms the
vector and magnitude is heading to
the direction of Cl caused a small
dipole moment present in molecule
hence polar
Covalent molecule
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7499
polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7699
whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8399
usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9399
example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9499
dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9899
water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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Diatomic molecule
Same element Different
Polyatomic molecule
As mmetric element
Non-polarmolecule
Polar molecule
al Symmetrical
Polarmolecule
Same surroundatoms Different surroundatoms
Non-polarmolecule
Polar molecule
326 Electronegativity and Type of Chemical Bond
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7399
more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7499
polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7699
whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7799
present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8099
Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9199
the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9399
example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9899
water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
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1 The type of bond that would form can be told by using the
difference of electronegativity (∆EN) Larger the difference the
more tendency of electron from low electronegativity atom to
move to the atom with higher electronegativity and form ioniccompound
a) The relationship between the ionic character and the difference
in the electronegativity of the bonded atom is shown in the
diagram and graph below
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b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
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more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
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polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
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density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
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present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
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983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
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charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
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c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8399
usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8799
are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8899
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9099
Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9199
the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
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Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9399
example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9499
dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
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Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
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below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
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Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9899
water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
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of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7299
b) From the graph above the dotted line represent the arbitrary line
between ionic and covalent characteristic of a molecule To be
more specific there more likely an ionic compound may have high
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7399
more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7499
polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7599
density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7699
whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7799
present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7899
983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7999
charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8099
Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8199
b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8299
c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8399
usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8499
3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8599
substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8699
34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8799
are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8899
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8999
3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9099
Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9199
the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9299
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9399
example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9499
dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9599
Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9699
below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9799
Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9899
water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9999
of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7399
more specific there more likely an ionic compound may have high
covalent characteristic (exemplified by LiI) or conversely covalent
compound having high ionic characteristic (exemplified by HF)
c) The covalent characteristic of a molecule is dependent on the ability
of a cation to polarise an anion Polarisation indicates the ability ofa cation to attract the electron density of an anion when put next to
the cation involved When a cation is able to pull the electron
electron with cation hence increase the covalency of the molecule
The covalency properties of a molecule is dependent on the cation
and anion where they can be explained qualitatively via
bull Polarisation power of cation bull Polarisability of anion
A+ 983128 991251
983106983083 983129983085
3611 Polarisation Power of Cation
Polarisation Power of Cation ndash measure the ability of a cation to
polarise the electron cloud of the anion
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7499
polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7599
density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7699
whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7799
present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7899
983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7999
charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8099
Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8199
b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8299
c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8399
usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8499
3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8599
substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8699
34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8799
are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8899
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8999
3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9099
Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9199
the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9299
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9399
example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9499
dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9599
Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9699
below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9799
Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9899
water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9999
of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7499
polarise the electron cloud of the anion
2 factors determining the polarisation power of cation
Charge of cation Size of cation
rArr Greater the charge of ion higher the
effective nuclear charge of cationhence it will be able to attract the
neighboring electron density of anion
rArr Smaller the size of cation closer the
neighboring anion to the nucleus ofcation hence easier for the cation to
polarise the anion and result an
This will caused the polarization power
of cation increase hence increase thecovalent characteristic of cation
increment in the polarization power of
cation and increase the covalentcharacteristic of cation
diams Both factors can be explained in another term called as charge density where
Charge Density = Charge Ionic Radius
diams From the equation above Charge Density will have a greater value provided that
cation has a high charge and small cationic radius
diams Greater the charge density higher the polarization power greater the covalent
characteristic of the cation
3612 Polarisability of Anion
bull Polarisability of an anion ~ ability of the anion to allow the electron
density to be polarised by cation
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7599
density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7699
whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7799
present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7899
983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7999
charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8099
Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8199
b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8299
c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8399
usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8499
3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8599
substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8699
34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8799
are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8899
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8999
3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9099
Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9199
the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9299
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9399
example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9499
dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9599
Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9699
below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9799
Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9899
water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9999
of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7599
density to be polarised by cation
bull 2 factors determining the polarisability of an anion
Charge of anion Size of anion
rArrGreater the charge of anion lower theeffective nuclear charge of anion This will
weakened the electrostatic attraction forces
rArrLarger the size of anion further theoutermost electron from the nucleus
of the anion easier for the cation to
bull Unlike cation anion does not have a term that combined bothfactors of charge and ionic radius However information of
polarisability of anion enable the prediction of the covalent
characteristic of a molecule since in order to form a covalent bond
it depend on both polarisation power of cation and polarisability ofthe anion
e ween nuc eus an e ou ermos
electron in anion and increase the polarisability of the anion hence increase
the covalent characteristic of anion
po ar se e an on an cause e
polarisability to increase henceincrease the covalent characteristic
of anion
362 Prediction of Chemical Bond Fajansrsquo Rule
bull In 1923 Kazimierz Fajans formulated an easy guidance to predict
whether a chemical bond will be covalent or ionic and depend on
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7699
whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7799
present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7899
983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7999
charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8099
Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8199
b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8299
c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8399
usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8499
3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8599
substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8699
34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8799
are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8899
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8999
3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9099
Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9199
the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9299
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9399
example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9499
dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9599
Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9699
below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9799
Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9899
water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9999
of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7699
whether a chemical bond will be covalent or ionic and depend on
the charge on the cation and the relative sizes of the cation and
anion They can be summarized in the following table
Ionic compound Low positive charge Large cation Small anion
Covalent compound High positive charge Small cation Large anion
bull Based on these guidance the bonding of a few compoundsshall be discussed to understand the application of Fajansrsquo
Rule in the chemical bonding
Lithium halide (LiX)
bull Lithium ion Li+ (1s2) has a small size due to only 1 shell
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7799
present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7899
983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7999
charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8099
Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8199
b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8299
c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8399
usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8499
3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8599
substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8699
34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8799
are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8899
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8999
3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9099
Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9199
the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9299
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9399
example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9499
dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9599
Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9699
below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9799
Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9899
water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9999
of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7799
present in its ion But since it has a low charge so its chargedensity is not too high That is why all lithium halide are ionic
compound The covalency of lithium halide varies from a
highly ioniccharacteristic to highly covalency depending on
the polarisability of the anion next to Li+
bull When a group of halide F ndash Cl ndash Br ndash I ndash is put close to Li+ the
Group 17 halide LiF is highly ionic since the fluoride ion hassmall ionic size and low charge hence has low polarisability
Ionic size increase with the increasing shell when going down
to Group 17 halide hence increase the polarisability which
allowed lithium ion to polarise the anionrsquos electron densityhence increase the covalency
983107983148 991251
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7899
983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7999
charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8099
Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8199
b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8299
c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8399
usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8499
3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8599
substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8699
34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8799
are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8899
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8999
3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9099
Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9199
the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9299
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9399
example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9499
dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9599
Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9699
below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9799
Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9899
water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9999
of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7899
983116983145983083
983110 991251
983106983154 991251
983107983148
Aluminium halide (AlX3) and aluminium oxide (Al2O3)
bull Aluminium ion (Al3+) has high charge density due to its high
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7999
charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8099
Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8199
b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8299
c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8399
usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8499
3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8599
substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8699
34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8799
are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8899
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8999
3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9099
Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9199
the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9299
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9399
example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9499
dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9599
Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9699
below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9799
Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9899
water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9999
of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 7999
charge unit and its small ionic radius So depending on the anionaluminium has a high tendency to form covalent compound For
example when going down to Group 17 halide aluminium fluoride
(AlF3) forms ionic compound (since F- has a low polarisability)
while aluminium trichloride (AlCl3) aluminium tribromide (AlBr 3)and aluminium iodide (AlI3) form covalent compound (since
chloride bromide and iodide have high polarisability) This
(10400
C) while aluminium trichloride and tribromide are 1920
C and780C respectively
bull As for aluminium oxide (Al2O3) it is an ionic compound with high
covalent characteristic as aluminium ion has high covalent
characteristic due to its high charge density This explained thehigh melting point of Al2O3 (20500C) yet it is insoluble in water It
also explained the amphoteric properties of aluminium oxide where
aluminium oxide can act as an acid (covalent characteristic) as
well as a base (ionic characteristic)
33 Metallic Bonding
1 Metallic bonding occurs when large
numbers of metal atoms interact
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8099
Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8199
b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8299
c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8399
usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8499
3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8599
substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8699
34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8799
are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8899
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8999
3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9099
Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9199
the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9299
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9399
example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9499
dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9599
Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9699
below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9799
Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9899
water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9999
of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8099
Unlike the reaction of metal with
non metal which involve electrons
transfer when two metal atoms
interact they can also share their
valence electrons in a covalent
bond and form gaseous diatomic
-
model of metallic bonding
proposes that all the metal atoms
in the sample contribute their
valence electrons to form an
electron ldquoseardquo that is
delocalized throughout the
piece The metal ions are
submerged within this electron
sea in an orderly array
a) The model we will use to study metallic bonding is band theory
because it states that delocalized electrons move freely through
ldquobandsrdquo formed by overlapping molecular orbitals
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8199
b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8299
c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8399
usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8499
3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8599
substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8699
34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8799
are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8899
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8999
3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9099
Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9199
the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9299
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9399
example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9499
dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9599
Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9699
below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9799
Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9899
water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9999
of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8199
b) Consider magnesium for example The electron configuration of
Mg is 1s22s22p63s2 so each atom has two valence electrons in the
3s orbital In a metallic crystal the atoms are packed closelytogether so the energy levels of each magnesium atom are
affected by the immediate neighbors of the atom as a result of
the energy scale that they are more appropriately described as aldquobandrdquo The closely spaced filled energy levels make up the
valence band The upper half of the energy levels corresponds
to the empty delocalized molecular orbitals formed by the overlap
of the 3p orbitals This set of closely spaced empty levels is calledthe conduction band As a conductor the conduction band and
valence band are overlapped hence electrons can travel freely
among the two bands hence conduct electricity
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8299
c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8399
usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8499
3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8599
substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8699
34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8799
are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8899
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8999
3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9099
Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9199
the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9299
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9399
example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9499
dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9599
Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9699
below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9799
Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9899
water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9999
of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8299
c) Theoretically greater the number of valence electrons in a
metal greater the number of electrons delocalised higher
the conductivity However the conductivity decrease withtemperature as vibration of the lattice of ion impedes the free
movement of electron in conduction band
2 Semiconductors are element that normally are not conductors but
will conduct electricity at elevated temperatures or when combined
with a small amount of certain other elements These elements are
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8399
usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8499
3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8599
substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8699
34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8799
are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8899
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8999
3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9099
Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9199
the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9299
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9399
example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9499
dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9599
Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9699
below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9799
Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9899
water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9999
of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8399
usually metalloid such as silicon and germanium(a) The energy gap between the conduction band and valence band
of these solids is much smaller than that for insulator If the energy
needed to excite electrons from the valence band into the conduction
band is provided the solid becomes a conductor Note that thisbehavior is opposite that of the metals
the conductivity of the semiconductors Doping can be done by
adding one of the following
i dopant atoms containing fewer valence electrons Hence the
semiconductor formed is positive p - type semiconductor
ii dopant atoms with extra valence electrons Hence the semiconductorformed is negative n - type semiconductor
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8499
3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8599
substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8699
34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8799
are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8899
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8999
3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9099
Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9199
the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9299
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9399
example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9499
dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9599
Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9699
below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9799
Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9899
water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9999
of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8499
3 Insulators are substances that do not conduct electricity
no matter how high the temperature is applied to the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8599
substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8699
34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8799
are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8899
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8999
3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9099
Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9199
the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9299
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9399
example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9499
dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9599
Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9699
below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9799
Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9899
water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9999
of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8599
substances involved The energy gaps between theconduction band and valence band of these element is very
large hence regardless how much energies were applied to
these insulator it will not be able to conduct electricity nor
heat Glass and woods are good examples of insulator Inwood and glass the gap between the valence band and the
Consequently much more energy is needed to excite anelectron into the conduction band Lacking this energy
electrons cannot move freely Therefore glass and wood
are insulators ineffective conductors of electricity
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8699
34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8799
are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8899
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8999
3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9099
Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9199
the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9299
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9399
example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9499
dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9599
Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9699
below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9799
Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9899
water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9999
of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8699
34 Intermolecular forces Van der Waals forces and hydrogen
bonding
1 The nature of the state of matter of substances and their changes
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8799
are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8899
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8999
3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9099
Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9199
the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9299
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9399
example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9499
dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9599
Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9699
below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9799
Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9899
water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9999
of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8799
are due primarily to forces among the molecules Both bonding
(intramolecular) forces and intermolecular forces arise from
electrostatic attractions between opposite charges Bonding forces
are due to the attraction between cations and anions (ionic
bonding) nuclei and electron pairs (covalent bonding) or metal
cations and delocalized valence electrons (metallic bonding)
between molecules as a result of partial charges or the attraction
between ions and molecules The two types of forces differ in
magnitude and forces explains why
a) Bonding forces are relatively strong because they involve
larger charges that are closer togetherb) Intermolecular forces are relatively weak because they typically
involve smaller charges that are farther apart
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8899
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8999
3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9099
Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9199
the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9299
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9399
example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9499
dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9599
Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9699
below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9799
Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9899
water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9999
of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8899
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8999
3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9099
Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9199
the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9299
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9399
example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9499
dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9599
Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9699
below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9799
Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9899
water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9999
of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 8999
3 Induced dipole Forces (Dispersion forces) - Consider how a helium
atom (monoatomic gas which have dipole moment = 0) interact
with the following species
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9099
Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9199
the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9299
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9399
example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9499
dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9599
Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9699
below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9799
Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9899
water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9999
of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9099
Helium with cationHelium with polar
moleculeHelium with Helium
He
He
a) From the diagram we can tell that if an ion or a polar
molecule is placed near an atom or a non-polar molecule the
electron distribution of the atom (or molecule) is distorted by
th f t d b th i th l l l lti i
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9199
the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9299
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9399
example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9499
dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9599
Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9699
below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9799
Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9899
water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9999
of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9199
the force exerted by the ion or the polar molecule resulting in
a kind of dipole The dipole in the atom (or non-polar
molecule) is said to be an induced dipole because the
separation of positive and negative charges in the atom(or non-polar molecule) is due to the proximity of an ion or a
polar molecule
b) However the weak attractive interaction between a non-polar
molecule to another non-polar molecule are unlike when
placed near an ion or polar molecule Between two non-polar
atom they form among themselves a short induced dipole
hence attract each other temporary Therefore the forcesformed between them are very weak and can be broken
easily Such interaction is also known as London forces
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9299
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9399
example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9499
dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9599
Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9699
below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9799
Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9899
water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9999
of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9299
Alkane CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18
RMM 16 30 44 58 72 86 100 114
Boiling ndash ndash ndash ndash
point oC
TrendRelative molecular mass increased
weak Van Der Waals forces increased
d) For molecules especially organic compounds which have the
same molecular mass and functioning group they may have
different boiling point depend on the molecular structure For
example pentane C5H12 with molecular mass 72 has 3 isomers
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9399
example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9499
dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9599
Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9699
below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9799
Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9899
water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9999
of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9399
example pentane C5H12 with molecular mass 72 has 3 isomersas shown in table below
Molecule 22-dimethylpropane 2-methylbutane n-pentane
structure
Boiling point
oC54 218 363
Total surface area increased
weak Van Der Waals forces increased
Boiling point increased
4 Dipole-Dipole Forces ~ When polar molecules lie near one
another as in liquids and solids their partial charges act as
tiny electric fields that orient them and give rise to dipole-
dipole forces the positive pole of one molecule attracts the
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9499
dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9599
Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9699
below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9799
Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9899
water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9999
of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9499
dipole forces the positive pole of one molecule attracts the
negative pole of another These are the forces that give polar
compound a higher boiling point than the non-polar
compound
Molecule FormulaMolecular
mass
Dipole
moment
Boiling point
(K)
Propane CH CH CH 44 0 08 231
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9599
Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9699
below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9799
Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9899
water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9999
of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9599
Propane CH3CH2CH3 44 008 231
Dimethyl
ether CH3OCH3 46 13 248
Methyl
chlorideCH3Cl 505 187 249
Ethanal CH3CHO 44 269 294
Acetonitrile CH3CN 41 392 355
Almost the same molecular massHowever greater the dipole moment
Stronger the dipole-dipole forces
Higher the boiling point
5 Hydrogen bond ~ a special type of dipole-dipole interaction
between the hydrogen atom in a polar bond as in NminusH OminusH
or FminusH with an electronegative O N or F atom Diagram
below shows a few example of interaction between molecules
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9699
below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9799
Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9899
water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9999
of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9699
below shows a few example of interaction between molecules
using hydrogen bond
Hydrogen bond
StF is more electronegative than N
Even though F is more electronegative than O amp
Hydrogen bond between FndashH is stronger than OndashH
However water form more hydrogen bond
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9799
Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9899
water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9999
of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9799
Strong
hydrogen
bond
F is more electronegative than NHydrogen bond between FndashH is stronger than NndashH
Higher the boiling point
b) The factors of hydrogen bonding can also use to explain the
solubility of some organic compound in water like example
ethane cannot dissolve in water but ethanol can dissolve in
water due to the hydrogen bonding
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9899
water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9999
of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9899
water due to the hydrogen bonding
c) Some organic compound form dimer using hydrogen bond
For example when glacial ethanoic acid is dissolved inorganic solvent it form a dimer using hydrogen bond via the
interaction between O-H and C=O in each of the molecule
d) In some case hydrogen bond can also be used to form which
is the intermolecular forces and intramolecular forces For
example in 2-nitrophenol and 4-nitrophenol the boiling point
of the 2 compounds can be explain below
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9999
of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol
8112019 Chemistry Pre-u Chemistry Sem 1 Chap 3
httpslidepdfcomreaderfullchemistry-pre-u-chemistry-sem-1-chap-3 9999
of the 2 compounds can be explain below
Since 2-nitrophenol form strong hydrogen bond asintramolecular forces the interaction between 2-nitrophenol
molecules are weaker among each other compare to 4-
nitrophenol which used hydrogen bond as theirintermolecular forces With stronger hydrogen bond which act
as the intermolecular forces the boiling point of 4-nitrophenol
is expected to be higher than 2-nitrophenol