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Transcript of Module 5 Chemical Bonds.doc
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MODULE 5
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Why do atoms bond together? - 'electron glue'!
Some atoms are very reluctant to combine with other atoms and exist in the air around
us as single atoms. These are the Noble Gases and have very stable electron
arrangements e.g. 2, 2,8 and 2,8,8 because their outer shells are full. The first three are
shown in the diagrams below and explains why Noble Gases are so reluctant to form
compounds with other elements.
(Atomic number) electron arrangement
All other atoms therefore, bond together to become electronically more stable, that is to
become like Noble Gases in electron arrangement. Bonding produces new substances and
usually involves only the 'outer shell' or 'valency' electrons and atoms can bond in two ways.
The phrase CHEMICAL BOND refers to the strong electrical force of attraction
between the atoms or ions in the structure. The combining power of an atom is
sometimes referred to as its valency and its value is linked to the number of outer
electrons of the original uncombined atom (see examples later).
(a) IONIC BONDING - By one atom transferring electrons to another atom to form
oppositely charged particles called ions which attract each other - the ionic bond.
An ion is an atom or group of atoms carrying an overall positive or negative
charge
o E.g. Na+, Cl-, [Cu(H2O)]2+, SO4
2- etc.
If a particle, as in a neutral atom, has equal numbers of protons (+) and electrons
(-) the particle charge is zero i.e. no overall electric charge.
The proton/atomic number in an atom does not change BUT the number of
associated electrons can!
If negative electrons are lost the excess charge from the protons produces an
overall positive ion.
If negative electrons are gained there is an excess of negative charge, so a negative
ion is formed.
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The charge on the ion is numerically related to the number of electrons transferred i.e.
electrons lost or gained.
For any atom or group of atoms, for every electron gained you get a one unit
increase in negative charge on the ion, for every electron lost you get a one unit
increase in the positive charge on the ion.
The atom losing electrons forms a positive ion (cation) and is usually a metal.
The atom gaining electrons forms a negative ion (anion) and is usually a non-
metallic element. The ionic bond then consists of the attractive force between the
positive and negative ions in the structure.
The ionic bonding forces act in all directions around a particular ion, it is not
directional, as in the case of covalent bonding.
(b) COVALENT BONDING - sharing electrons to form molecules with covalent bonds,
the bond is usually formed between two non-metallic elements in a molecule. The two
positive nuclei (due to the positive protons in them) of both atoms are mutually
attracted to the shared negative electrons between them - the covalent bond. They share
the electrons in a way that gives a stable Noble Gas electron arrangement.
This kind of bond or electronic linkage does act in a particular direction i.e. along
the 'line' between the two nuclei of the atoms bonded together; this is why molecules
have a particular shape.
(c) METALLIC BONDING isn't quite like ionic or covalent bonding, the metal atoms
form positive ions, but no negative ion is formed from the same metal atoms, but the
positive metal ions/atoms are attracted together by the free moving negative electronsbetween them.
NOBLE GASES are very reluctant to share, gain or lose electrons to form a chemical
bond. They are already electronically very stable. For most other elements the types of
bonding and the resulting properties of the elements or compounds are described in
detail in Parts 2 to 5. In all the electronic diagrams ONLY the outer electrons are
shown.
IONIC BONDING - compounds and properties
Examples of ionic compounds*physical properties of ionic compounds
Ionic Bonding - electron transfer
Ionic bonds are formed by one atom transferring electrons to another atom to form ions.
Elements consist of neutral atoms or molecules, the electrical neutrality is because thenumber of positive protons equals the number of negative electrons .
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Ions are atoms, or groups of atoms, which have lost or gained electrons to have a net
electrical charge overall.
The atom losing electrons forms a positive ion (a cation) and is usually a metal. The overall
charge on the ion is positive due to excess positive nuclear charge (protons do NOT change in
chemical reactions).
The atom gaining electrons forms a negative ion (an anion) and is usually a non-metallic
element. The overall charge on the ion is negative because of the gain, and therefore excess,
of negative electrons.
The examples below combining a metal from Groups 1 (Alkali Metals), 2 or 3, with a non-
metal from Group 6 or Group 7 (The Halogens). The electron structures are shown in () or [].
Only the outer electrons of the original atoms, and where they end up in the ions, are shown
in the dot and cross (ox) diagrams
Ionic bonding is not directional like covalent bonding, in the sense that the force of attractionbetween the positive ions and the negative ions act in every direction around the ions.
Example 1: A Group 1 metal + a Group 7 non-metal e.g. sodium + chlorine ==> sodium
chloride NaCl or ionic formula Na+Cl- In terms of electron arrangement, the sodium donates
its outer electron to a chlorine atom forming a single positive sodium ion and a single
negative chloride ion. The atoms have become stable ions, because electronically, sodium
becomes like neon and chlorine like argon.
Na (2.8.1) + Cl (2.8.7) ==> Na+
(2.8) Cl-
(2.8.8)
can be summarised electronically to give the stable 'noble gas' structures as [2, 8,1] + [2,8,7]
==> [2,8]+ [2,8,8]-
ONE combines with ONE to form
The valencies of Na and Cl are both 1, that is, the numerical charge on the ions. sodiumfluoride NaF, potassium bromide KBr and lithium iodide LiI etc. will all be electronically
similar.
Note:
The charge on the sodium ion Na+ is +1 units (shown as just +) because there is one more
positive proton than there are negative electrons in the sodium ion.
The charge on the chloride ion Cl- is -1 unit (shown as just -) because there is one more
negative electron than there are positive protons in the chloride ion.
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Example 2: A Group 2 metal + a Group 7 non-metal e.g. magnesium + chlorine ==>
magnesium chloride MgCl2 or ionic formula Mg2+(Cl-)2 In terms of electron arrangement,
the magnesium donates its two outer electrons to two chlorine atoms forming a double
positive magnesium ion and two single negative chloride ions. The atoms have become stable
ions, because electronically, magnesium becomes like neon and chlorine like argon.
Mg (2.8.2) + 2Cl (2.8.7) ==> Mg2+ (2.8) 2Cl- (2.8.8)
can be summarised electronically as [2,8,2] + 2[2,8,7] ==> [2,8]2+ [2,8,8]-2
ONE combines with TWO to form see *
* NOTE you can draw two separate chloride ions, but in these examples square brackets and
a number subscript have been used, as in ordinary chemical formula. The valency of Mg is 2
and chlorine 1, i.e. the numerical charges of the ions. Beryllium fluoride BeF 2, magnesium
bromide MgBr2, calcium chloride CaCl2 or barium iodide BaI2 etc. will all be electronically
similar.
Ca is 2.8.8.2, F is 2.7 rest of dot and cross diagrams are up to you.
Example 3: A Group 3 metal + a Group 7 non-metal e.g. aluminium + fluorine ==>
aluminium fluoride AlF3 or ionic formula Al3+(F-)3 In terms of electron arrangement, thealuminium donates its three outer electrons to three fluorine atoms forming a triple positive
aluminium ion and three single negative fluoride ions. The atoms have become stable ions,
because aluminium and fluorine becomes electronically like neon. Valency of Al is 3 and F is
1, i.e. equal to the charges on the ions.
Al (2.8.3) + 3F (2.7) ==> Al3+ (2.8) 3F- (2.8)
can be summarised electronically as [2,8,3] + 3[2,7] ==> [2,8]3+ [2,8]-3
ONE combines with THREE to form
Solid aluminium chloride/bromide/iodide have similar formula but are covalent when
vapourised into Al2X6 dimer molecules.
Example 4: A Group 1 metal + a Group 6 non-metal e.g. sodium/potassium + oxygen ==>
sodium/potassium oxide Na2O/K2O or ionic formula (Na+)2O2-/(K+)2O2- In terms of
electron arrangement, the two sodium/potassium atoms donate their outer electron to one
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oxygen atom. This results in two single positive potassium ions to one double negative oxide
ion. All the ions have the stable electronic structures 2.8.8 (argon like) or 2.8 (neon like).
Valencies, K 1, oxygen 2. Lithium oxide, Li2O, sodium oxide Na2O, sodium sulphide Na2S
and potassium K2S etc. will be similar.
Sodium oxide
2Na (2.8.1) + O (2.6) ==> 2Na+ (2.8.8) O2- (2.8)
can be summarised electronically as 2[2,8,1] + [2,6] ==> [2,8]+2 [2,8]2-
TWO combine with ONE to form
or
+ ==>
Potassium oxide
2K(2.8.8.1) + O (2.6) ==> 2K+ (2.8.8) O2- (2.8)
can be summarised electronically as 2[2,8,8,1] + [2,6] ==> [2,8,8]+2 [2,8]2-
TWO combine with ONE to form
The electronic similarities between the two examples are very obvious.
Li is 2.1, Na is 2.8.1, S is 2.8.6 (for group 1 sulphide compound), rest of dots and crosses
diagrams are up to you.
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Example 5: A Group 2 metal + a Group 6 non-metal e.g. magnesium/calcium + oxygen
==> magnesium/calcium oxide MgO/CaO or ionic formula Mg2+O2-/Ca2+O2- In terms of
electron arrangement, one magnesium/calcium atom donates its two outer electrons to one
oxygen atom. This results in a double positive calcium ion to one double negative oxide ion.
All the ions have the stable electronic structures 2.8.8 (argon like) or 2.8 (neon like). The
valency of both calcium and oxygen is 2.
Magnesium oxide
ONE combines with ONE to form
For magnesium oxide: Mg (2.8.2) + O (2.6) ==> Mg2+ (2.8) O2- (2.8)
The stable 'noble gas' structures can be summarised electronically as [2,8,2] + [2,6] ==>
[2,8,8]2+ [2,8]2-
Calcium oxide
Ca (2.8.8.2) + O (2.6) ==> Ca2+ (2.8.8) O2- (2.8)
can be summarised electronically as [2,8,8,2] + [2,6] ==> [2,8,8]2+
[2,8]2-
ONE combines with ONE to form
Magnesium oxide MgO, magnesium sulphide MgS and calcium sulphide CaS will be similar
electronically and give identical giant ionic lattice structures. Group 2 metals lose the two
outer electrons to give the stable 2+ positive ion (cation) and S and O, both non-metals in
Group 6, have 6 outer electrons and gain 2 electrons to form 2- negative ion (anion).
Formagnesium sulphide: Mg (2.8.2) + S (2.8.6) ==> Mg2+ (2.8) S2- (2.8.8)
Forcalcium sulphide: Ca (2.8.8.2) + S (2.8.6) ==> Ca2+ (2.8.8) S2- (2.8.8)
The dot and cross (ox) diagrams will be identical to that for calcium oxide above, except Mg
instead of Ca (same group) and S instead of O (same group of Periodic Table).
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Example 6: A Group 3 metal + a Group 6 non-metal e.g. aluminium + oxygen ==>
aluminium oxide Al2O3 or ionic formula (Al3+)2(O
2-)3 In terms of electron arrangement, two
aluminium atoms donate their three outer electrons to three oxygen atoms. This results in two
triple positive aluminium ions to three double negative oxide ions. All the ions have the
stable electronic structure of neon 2.8. Valencies, Al 3 and O 2.
2Al (2.8.3) + 3O (2.6) ==> 2Al3+ (2.8) 3O2- (2.8)
can be summarised electronically as 2[2,8,3] + 3[2,6] ==> [2,8]3+2 [2,8]2-
3
TWO combine with THREE to form
Note:
The charge on the aluminium ion Al3+ is +3 units (shown as 3+) because there are three more
positive protons than there are negative electrons in the aluminium ion.
The charge on the oxide ion O2- is -2 units (shown as 2-) because there are two more negative
electrons than there are positive protons in the oxide ion.
The properties of Ionic Compounds
The diagram on the right is typical of the
giant ionic crystal structure of ionic
compounds like sodium chloride and
magnesium oxide.
The alternate positive and negative ions in
an ionic solid are arranged in an orderlyway in a giant ionic lattice structure shown
on the left.
The ionic bond is the strong electrical attraction between the positive
and negative ions next to each other in the lattice.
The bonding extends throughout the crystal in all directions.
Salts and metal oxides are typical ionic compounds.
This strong bonding force makes the structure hard (if brittle) and has
high melting and boiling points, so they are not very volatile!
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A relatively large amount of energy is needed to melt or boil ionic
compounds.
The bigger the charges on the ions the stronger the bonding attraction
e.g. magnesium oxide Mg2+O2- has a higher melting point than sodium
chloride Na+
Cl-
.
Unlike covalent molecules, ALL ionic compounds are crystalline solids
at room temperature.
They are hard but brittle, when stressed the bonds are broken along
planes of ions which shear away. They are NOT malleable like metals
(see below).
Many ionic compounds are soluble in water, but not all, so don't make
assumptions. Salts can dissolve in water because the ions can separate
and become surrounded by water molecules which weakly bond to the
ions. This reduces the attractive forces between the ions, preventing thecrystal structure to exist. Evaporating the water from a salt solution will
eventually allow the ionic crystal lattice to reform.
The solid crystals DO NOT conduct electricity because the ions are not
free to move to carry an electric current. However, if the ionic
compound is melted or dissolved in water, the liquid will now conduct
electricity, as the ion particles are now free.
3. Covalent Bonding - electron sharing in big or small molecules!
Covalent bonds are formed by atoms sharing electrons to form molecules. This type of
bond usually formed between two non-metallic elements. The molecules might be that of
an element i.e. one type of atom only OR from different elements chemically combined
to form a compound.
The covalent bonding is caused by the mutual electrical attraction between the twopositive nuclei of the two atoms of the bond, and the negative electrons between them.
One single covalent bond is a sharing of 1 pair of electrons, two pairs of shared
electrons between the same two atoms gives a double bond and it is possible for two
atoms to share 3 pairs of electrons and give a triple bond.
Note: In the examples it is assumed you can work out the electron configuration
(arrangement in shells or energy levels) given the atomic number from the Periodic
Table.
This kind of bond or electronic linkage does act in a particular direction i.e. along the'line' between the two nuclei of the atoms bonded together; this is why molecules have a
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particular shape. In the case of ionic or metallic bonding, the electrical attractive forces act
in all directions around the particles involved.
The bonding in Small Covalent Molecules
The simplest molecules are formed from two atoms and examples of their formation are
shown below. The electrons are shown as dots and crosses to indicate which atom the
electrons come from, though all electrons are the same. The diagrams may only show
the outer electron arrangements for atoms that use two or more electron shells. The
electron structures are given in (). Examples of simple covalent molecules are
Example 1: two hydrogen atoms (1) form the molecule of the element hydrogen H2
and combine to form where both atoms have a pseudo helium
structure of 2 outer electrons around each atom's nucleus. Any covalent bond is formed from
the mutual attraction of two positive nuclei and negative electrons between them. The nuclei
would be a tiny dot in the middle of where the H symbols are drawn! H valency is 1.
Example 2: two chlorine atoms (2.8.7) form the molecule of the element chlorine Cl2
and combine to form where both atoms have a
pseudo argon structure of 8 outer electrons around each atom. All the other halogens would
be similar e.g. F2, Br2 and I2 etc. Valency of halogens like chlorine is 1 here.
Example 3: one atom of hydrogen (1) combines with one atom of chlorine (2.8.7) to form the
molecule of the compound hydrogen chloride HCl
and combine to form where hydrogen is electronically like
helium and chlorine like argon. All the other hydrogen halides will be similar e.g. hydrogen
fluoride HF, hydrogen bromide HBr and hydrogen iodide HI.
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Note: Hydrogen chloride gas is a true covalent substance consisting of small HCl molecules.
If the gas is dissolved in a hydrocarbon solvent like hexane or methylbenzene it remains as
HCl molecules and because there are no ions present, the solution does not conduct
electricity. However, if hydrogen chloride gas is dissolved in water, things are very different
and the HCl molecules split into ions. Hydrochloric acid is formed which consists of a
solution ofhydrogen ions (H+) and chloride ions (Cl-). The solution then conductselectricity and passage of a d.c. current causes electrolysis to take place forming hydrogen
and chlorine.
Example 4: two atoms of hydrogen (1) combine with one atom of oxygen (2.6) to form the molecule
of the compound water H2O
and and combine to form so that the hydrogen atoms are
electronically like helium and the oxygen atom becomes like neon. The molecule can be shown as
with two hydrogen - oxygen single covalent bonds (AS note: called a V or bent shape, the
H-O-H bond angle is 105o). Hydrogen sulphide will be similar, since sulphur (2.8.6) is in the same
Group 6 as oxygen. Valency of oxygen and sulphur is 2 here.
Example 5: three atoms of hydrogen (1) combine with one atom of nitrogen (2.5) to form themolecule of the compound ammonia NH3
three of and one combine to form so that the hydrogen
atoms are electronically like helium and the nitrogen atom becomes like neon. The molecule
can be shown
as with three nitrogen - hydrogen single covalent bonds (AS note: called a trigonal pyramid
shape, the H-N-H bond angle is 107o). PH3 will be similar since phosphorus (2.8.5) is in the
same Group 5 as nitrogen. Valency of nitrogen or phosphorus is 3 here.
Example 6: four atoms of hydrogen (1) combine with one atom of carbon (2.4) to form themolecule of the compound methane CH4
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four of and one of combine to form so that the hydrogen
atoms are electronically like helium and the nitrogen atom becomes like neon. The molecule
can beshown as with four carbon - hydrogen single covalent bonds (AS note:
called a tetrahedral shape, the H-C-H bond angle is 109o). SiH4 will be similar because
silicon (2.8.4) is in the same group as carbon.
All the bonds in the above examples are single covalent bonds. Below are three examples 7-
9, where there is a double bond in the molecule, in order that the atoms have stable Noble
Gas outer electron arrangements around each atom. Carbon and silicon have a valency of 4.
More complex examples can be worked out e.g. involving C, H and O. In each case link in
the atoms so that there are 2 around a H (electronically like He), or 8 around the C or O
(electronically like Ne).
Example 7: Two atoms of oxygen (2.6) combine to form the molecules of
the element oxygen O2. The molecule has one O=O double covalent bond .
Oxygen valency 2.
Example 8: One atom of carbon (2.4) combines with two
atoms of oxygen (2.6) to form the compound carbon dioxide CO2. The molecule can be
shown as with two carbon = oxygen double covalent bonds (AS note:
called a linear shape, the O=C=O bond angle is 180o). Valencies of C and O are 4 and 2
respectively.
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Example 9: Two atoms of carbon (2.4) combine with four atoms of
hydrogen (1) to form ethene C2H4. The molecule can be shown as with one
carbon = carbon double bond and four carbon - hydrogen single covalent bonds (called a planar
shape, its completely flat!, the H-C=C and H-C-H bond angles are 120o). The valency of carbon is still
4.
Examples 10-13: The scribbles below illustrate some more complex examples. Can you
deduce them for yourself? Ex. 10nitrogen N2; Ex. 11ethane C2H6; Ex. 12chloromethane
CH3Cl and Ex. 13 methanol CH3OH. Electronic origin of the diagrams showing the outer
electrons of N, C, Cl and O: N at. no. 7 (2.5), H at. no. (1), C at. no. 6 (2.4), Cl at. no. 17
(2.8.7) and O at. no. 8 (2.6) plus a variety of crosses and blobs! The valencies or combining
power in these examples are N 3, H 1, C 4, Cl 1 and O 2. From these you can work out others
e.g. Ex. 12 can be used to derive the ox diagram for tetra-chloromethane CCl4.
AS advanced level notes on shapes and bond angles:
o Ex. 11 Ethane has a linked double tetrahedral shape, all H-C-H and H-C-C
bond angles are 109o
o Ex. 12 chloromethane has tetrahedral shape with H-C-H and H-C-Cl bond
angles of approximately 109o
o Ex. 13 methanol, the four bonds around the central carbon are tetrahedrally
arranged with a H 'wiggle' on the oxygen. All the H-C-H, H-C-O and C-O-H
bond angles are approximately 109o
o The blue icon e.g. below, represents an octahedral shape (e.g. SF6, complex
transition metal ions like [Cu(H2O)6]2+ and the bond angles are either 90o or
180o
o Simple molecules with a triple bond are often linear e.g. H-C C-H ethyne
orH-C N hydrogen cyanide (methanenitrile)
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Typical properties of simple covalent substances - small molecules!
Covalent substances tend to be liquids or gases at room temperature because the forces
between their particles (the molecules) are weak. These intermolecular forces must not be
confused with intramolecular forces. Inter- means between; intra- means within.
Intermolecular forces are forces between individual molecules (such as van der Waal
forces). They are weak forces.
Intramolecular forces are forces within the molecules (covalent bonds). They hold the
atoms together in the molecule. These are strong forces.
-
+ +
-
Intra-molecular force +
Intermolecular force (between one water molecule and another)
-
+ +
Intermolecular forces
We will discuss two types of intermolecular forces:
van der Waals forces
hydrogen bonds
Van der Waals forces
All covalent molecules, whether polar or non polar, develop temporary or instantaneous
dipoles. This results from uneven movement of all the electrons within the molecules. Van
der waals forces are the weak attraction between oppositely charged ends of molecules with
temporary dipoles.
H
s
s
s
s
H
O
H H
O
H
H
O
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+ - + - + -
a molecule a molecule ..but can develop one
with a temporary that does
dipole have a dipole
Figure 4.0 Van der Waals forces. When molecule 2 develops a temporary dipole then it will be attracted
to molecule 1.
HYDROGEN BONDS
The hydrogen bond is the weak attraction between the slightly positive hydrogen atom in one
polar molecule and the slightly electronegative atom in another polar molecule of the same
type or of a different type.
Structure and Properties of simple covalent substances
The electrical forces of attraction that is the chemical bond, between atoms in a
molecule are usually very strong, so, most covalent molecules do not change
chemically on moderate heating.
o e.g. although a covalent molecule like iodine, I2, is readily vapourised on
heating, it does NOT break up into iodine atoms I. The I-I covalent bond is
strong enough to withstand the heating and the purple vapour still consists of
the same I2 molecules as the dark coloured solid is made up of.
So why the ease of vaporisation on heating?
o The electrical attractive forces between individual molecules are weak, so
the bulk material is not very strong physically and there are also
consequences for the melting and boiling points.
These weak electrical attractions are known as intermolecular forces and are
readily weakened further on heating. The effect of absorbing heat energy results in
increased the thermal vibration of the molecules which weakens the intermolecular
forces. In liquids the increase in the average particle kinetic energy makes it easier for
molecules to overcome the intermolecular forces and change into a gas or vapour.
Consequently, small covalent molecules tend to be volatile liquids with low boiling
points, so easily vapourised or low melting point solids.
o On heating the inter-molecular forces are easily overcome with the increased
kinetic energy of the particles giving the material a low melting orboiling
point and a relatively small amount of energy is needed to effect these state
changes.
o This contrasts with the high melting points of giant covalent structures with
their strong 3D network.
o Note: The weak electrical attractive forces between molecules, the so called
intermolecular forces should be clearly distinguished between the strong
covalent bonding between atoms in molecules (small or giant), and these aresometimes referred to as intramolecular forces (i.e. internal to the molecule).
12
12
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Covalent structures are usually poor conductors of electricity because there are no
free electrons or ions in any state to carry electric charge.
Most small molecules will dissolve in some solvent to form a solution.
o This again contrasts with giant covalent structures where the strong bond
network stops solvent molecules interacting with the particles making up the
material.
The properties of these simple small molecules should be compared and
contrasted with those molecules of a giant covalent nature (next section).
o Apart from points on the strong bonds between the atoms in the molecule and
the lack of electrical conduction, all the other properties are significantly
different!
Large Covalent Molecules and their Properties
(Macromolecules - giant covalent networks and polymers)
Because covalent bonds act in a particular direction i.e. along the 'line' between the two
nuclei of the atoms bonded together in an individual bond, strong structures can be
formed, especially if the covalent bonds are arranged in a strong three dimensional
giant covalent lattice.
The structure of the three allotropes of carbon (diamond,graphite and fullerenes), silicon and silicon dioxide (silica)
DIAGRAMS
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It is possible formany atoms to link up to form a giant
covalent structure or lattice. The atoms are usually non-
metals.
This produces a very strong 3-dimensional covalent
bond network or lattice.
This gives them significantly different properties from
the small simple covalent molecules mentioned above.
This is illustrated by carbon in the form of diamond
(an allotrope of carbon). Carbon has four outer electrons
that form four single bonds, so each carbon bonds to four
others by electron pairing/sharing. Pure silicon, another
element in Group 4, has a similar structure.
o NOTE: Allotropes are different forms of the same
element in the same physical state. They occur due
to different bonding arrangements and so
diamond, graphite (below) and fullerenes
(below) are the three solid allotropes of the
element carbon.
Oxygen (dioxygen), O2, and ozone
(trioxygen), O3, are the two small gaseous
allotrope molecules of the element oxygen.
Sulphur has three solid allotropes, two
different crystalline forms based on smallS8 molecules called rhombic and
monoclinic sulphur and a 3rd form of long
chain ( -S-S-S- etc.) molecules called
plastic sulphur.
TYPICAL PROPERTIES of GIANT COVALENT
STRUCTURES
This type of giant covalent structure is thermally very
stable and has a very high melting and boiling points
because of the strong covalent bond network (3D or 2D inthe case of graphite below).
A relatively large amount of energy is needed to melt or
boil giant covalent structures.
They are usually poor conductors of electricity because
the electrons are not usually free to move as they can in
metallic structures.
Also because of the strength of the bonding in all
directions in the structure, they are often very hard,
strong and will not dissolve in solvents like water. The
DIAMOND
SILICA
silicon dioxide
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bonding network is too strong to allow the atoms to
become surrounded by solvent molecules
Silicon dioxide (silica, SiO2) has a similar 3D structure
and properties to carbon (diamond) shown below.
The hardness of diamond enables it to be used as the
'leading edge' on cutting tools.
Many naturally occurring minerals are based on -O-
X-O- linked 3D structures where X is often silicon (Si)
and aluminium (Al), three of the most abundant
elements in the earth's crust.
o Silicon dioxide is found as quartz in granite
(igneous rock) and is the main component in
sandstone - which is a sedimentary rock, formed
the compressed erosion products of igneous rocks.
o Many some minerals that are hard wearing, rare
and attractive when polished hold great value as
gemstones.
Carbon also occurs in the form of graphite. The carbon
atoms form joined hexagonal rings forming layers 1 atom
thick.
There are three strong covalent bonds per carbon (3
C-C bonds in a planar arrangement from 3 of its 4 outerelectrons), BUT, the fourth outer electron is 'delocalised'
or shared between the carbon atoms to form the
equivalent of a 4th bond per carbon atom (this situation
requires advanced level concepts to fully explain, and this
bonding situation also occurs in fullerenes described
below, and in aromatic compounds you deal with at
advanced level).
The layers are only held together by weak
intermolecular forces shown by the dotted lines NOT by
strong covalent bonds.
Like diamond and silica (above) the large molecules of
the layer ensure graphite has typically very high melting
point because of the strong 2D bonding network(note:
NOT 3D network)..
Graphite will not dissolve in solvents because of the
strong bonding
BUT there are two crucial differences compared to
diamond ...
o Electrons, from the 'shared bond', can move
GRAPHITE
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freely through each layer, so graphite is a
conductor like a metal (diamond is an electrical
insulator and a poor heat conductor). Graphite is
used in electrical contacts e.g. electrodes in
electrolysis.
o The weak forces enable the layers to slip over
each other so where as diamond is hard material
graphite is a 'soft' crystal, it feels slippery.
Graphite is used as a lubricant.
These two different characteristics described above are
put to a common use with the electrical contacts in
electric motors and dynamos. These contacts (called
brushes) are made of graphite sprung onto the spinning
brass contacts of the armature. The graphite brushes
provide good electrical contact and are self-lubricating asthe carbon layers slide over each other.
A 3rd form of carbon are fullerenes or 'bucky balls'! It
consists of hexagonal rings like graphite and alternating
pentagonal rings to allow curvature of the surface.
Buckminster Fullerene C60 is shown and the bonds form
a pattern like a soccer ball. Others are oval shaped like a
rugby ball. It is a black solid insoluble in water.
They are NOT considered giant covalent structures
and are classed as simple molecules. They do dissolve in
organic solvents giving coloured solutions (e.g. deep red
in petrol hydrocarbons, and although solid, their melting
points are not that high.
They are mentioned here to illustrate the different
forms of carbon AND they can be made into continuous
tubes to form very strong fibres of 'pipe like' molecules
called 'nanotubes'. These 'molecular size' particles
behave quite differently to a bulk carbon material like
graphite.
Uses of Nanotubes:
o They can be used as semiconductors in electrical
circuits.
o They act as a component of industrial catalysts
for certain reactions whose economic efficiency is
of great importance (time = money in business!).
The catalyst can be attached to the
nanotubes which have a huge surface are
FULLERENES
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per mass of catalyst 'bed'.
They large surface combined with the
catalyst ensure two rates of reaction factors
work in harmony to increase the speed of
the industrial reaction.
o Nanotube fibres are very strong and so they are
used in 'composite materials' e.g. reinforcing
graphite in carbon fibre tennis rackets.
o Nanotubes can 'cage' other molecules and can be
used as a means of delivering drugs in controlled
way to the body.
BONDING IN METALS
METALLIC BONDING - structure and properties of metals
The crystal lattice of metals consists of ions NOTatomssurrounded by a 'sea of electrons' forming
another type ofgiant lattice.
The outer electrons(-) from the original metal atoms are free to move
around between the positive metal ions formed (+).
These free or 'delocalised' electrons are the 'electronic glue' holding the
particles together.
There is a strong electrical force of attractionbetween these mobile
electrons (-) and the 'immobile' positive metal ions (+) and this is themetallic bond.
Metallic bonding is not directional like covalent bonding, it is like ionic
bonding in the sense that the force of attraction between the positive metal
ions and the mobile electrons acts in every direction about the fixed
(immobile) metal ions.
Explaining the physical properties of metals
This strong bonding generally results in dense, strong materials with high melting
and boiling points. Usually a relatively large amount of energy is needed to melt or boil metals. .
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Metals are good conductors of electricitybecause these 'free' electrons carry the
charge of an electric current when a potential difference (voltage!) is applied across a
piece of metal.
Metals are also good conductors of heat. This is also due to the free moving
electrons. Non-metallic solids conduct heat energy by hotter more strongly vibratingatoms, knocking against cooler less strongly vibrating atoms to pass the particle
kinetic energy on. In metals, as well as this effect, the 'hot' high kinetic energy
electrons move around freely to transfer the particle kinetic energy more efficiently to
'cooler' atoms.
Typical metals also have a silvery surface but remember this may be easily tarnished
by corrosive oxidation in air and water.
Unlike ionic solids, metals are very malleable, they can be readily bent, pressed or
hammered into shape. The layers of atoms can slide over each other without
fracturing the structure (see below). The reason for this is the mobility of the
electrons. When planes of metal atoms are 'bent' or slide the electrons can run in
between the atoms and maintain a strong bonding situation. This can't happen in ionic
solids.
Note on Alloy Structure
1. Shows the regular arrangement of the atoms in a metal crystal and the white
spaces show where the free electrons are (yellow circles actually positive
metal ions).
2. Shows what happens when the metal is stressed by a strong force. The layers
of atoms can slide over each other and the bonding is maintained as the
mobile electrons keep in contact with atoms, so the metal remains intact BUT
a different shape.
3. Shows an alloy mixture. It is NOT a compound but a physical mixing of a
metal plus at least one other material (shown by red circle, it can be another
metal e.g. Ni, a non-metal e.g. C or a compound of carbon or manganese, and
it can be bigger or smaller than iron atoms). Many alloys are produced to give
a stronger metal. The presence of the other atoms (smaller or bigger) disrupts
the symmetry of the layers and reduces the 'slip ability' of one layer next to
another. The result is a stronger harder less malleable metal.
4. The main point about using alloys is that you can make up, and try out, all
sorts of different compositions until you find the one that best suits the
required purpose.
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POLAR AND NON-POLAR COVALENT COMPOUNDS
POLAR MOLECULES
A covalent bond in which the electron pair is shared unequally has partial ionic character and
is called a polar covalent bond. ( A polar covalent is a bond between two atoms that
have partial electric charges arising from their difference in electro-negativity. The
partial charges give rise to an electric dipole moment.)The extent to which an atom has a
greater or lesser share of electrons in a bond is determined by the difference in electro-
negativities of the two bonds. The electro-negativity of an element is its electron-pulling
power when it is part of a compound. An atom with a high electro-negativity has a strong
pulling power on electrons particularly for electron pair it shares with its neighbour. The
outcome of the tug-of-war: the more electronegative atom has a greater share of electron pair
of the covalent bond.
An O-H bond is polar because oxygen is more electronegative than hydrogen and gains a
greater share in the bonding electron pair. Its greater share of electrons means that oxygen
has a partial negative charge, which we denote -. Because the electron pair has been
pulled away from hydrogen atom, that atom has a partial positive charge, denoted +. We
show the partial charges on the atoms by writing +H-O - . A polar molecule is a molecule
with a nonzero dipole moment. All diatomic molecules composed of atoms of different
elements are slightly polar.
Examples of polar molecules
HF
HCl
HBr
HI - +
CO
ClF
H2O
NH3
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NON-POLAR MOLECULES
A non-polar molecule is molecule that has zero electric dipole moment. All homo-nuclear
(same) diatomic molecules, such as Cl2 and H2, are non-polar because there are no partial
charges on their atoms.
Examples of Non- polar Molecules
Cl2
H2
F2
N3
O2