Chemical Bonding and Molecular Structure
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Transcript of Chemical Bonding and Molecular Structure
Bonding and Structure 1
Chemical Bonding and
Molecular Structure
• Ionic vs. covalent bonding• Molecular orbitals and the covalent bond • Valence electron Lewis dot structures
octet vs. non-octetresonance structuresformal charges
• VSEPR (Valence shell electron pair repulsion) - predicting shapes of molecules• Bond properties
polarity, bond order, bond strength
Bonding and Structure 2
Chemical Bonding
Problems and questions —• How is a molecule or polyatomic
ion held together?• Why are atoms distributed at
strange angles?• Why are molecules not flat?• Can we predict the structure?• How is structure related to
chemical and physical properties?
Bonding and Structure 3
Most bonds are somewhere in between.
Forms of Chemical Bonds• There are 2 extreme forms of connecting
or bonding atoms:
• Ionic—complete transfer of electrons from one atom to another
• Covalent—electrons shared between atoms
Bonding and Structure 4
Ionic Ionic BondsBonds
Ionic compounds - essentially complete electron
transfer from an element of low (metal) to an element of high electron affinity (EA) (nonmetal)
Na(s) + 1/2 Cl2(g) Na+ + Cl-
NaCl (s)
- NON-DIRECTIONAL bonding via Coulomb (charge) interaction
- primarily between metals (Groups 1A, 2A and transition metals) and nonmetals (esp O and halogens)
Bonding and Structure 5
Covalent Bonding
Covalent bond is the sharing of the VALENCE ELECTRONS of each atom in a bond
Recall: Electrons are divided between core and valence electrons. ATOM core valenceNa 1s2 2s2 2p6 3s1 [Ne] 3s1
Br [Ar] 3d10 4s2 4p5 [Ar] 3d10 4s2 4p5
Br Br
Bonding and Structure 6
Valence Electrons1A1A
2A2A 3A3A 4A4A 5A5A 6A6A 7A7A
8A8A
Number of valence electrons is equal to the Group number.
Bonding and Structure 7
Covalent BondingThe bond arises from the mutual attraction of
2 nuclei for the same electrons.
HB+ HA HBHA
A covalent bond is a balanceof attractive and repulsive forces.
Bonding and Structure 8
←Fig. (a) In the nonpolar covalent bond present, there is a symmetrical distribution of electron density. (b) In the polar covalent bond present, electron density is displaced because of its electronegativity.
Polar and non-polar covalent bond
Dipole moment, µ= e (esu) x d (angstrom)
Greater the DM greater the polarity
Bonding and Structure 9
Bond FormationA bond can result from a “head-to-head/ end-to-end”
overlap of atomic orbitals on neighboring atoms.
H H Cl••
••
•• Cl
••••
••+
Overlap of H (1s) and Cl (2p)
This type of overlap places bonding electrons in a
MOLECULAR ORBITAL along the line between the two atoms and forms a SIGMA BOND ().
S-s, s-p and p-p orbitals form sigma bond.
Bonding and Structure 10
Sigma Bond Formation by Orbital Overlap
•• ••
sigma bond ( )
+HH
Two s Atomic Orbitals (A.O.s) overlap to form an s (sigma) Molecular Orbital (M.O.)
6_H2pot.mov
Bonding and Structure 11
Sigma Bond Formation by Orbital Overlap
•• ••
sigma bond ( )
+HH
Two s A.O.s overlap to from an s M.O.
Similarly, two p A.O.s can overlap end-on to from a p M.O.
e.g.F2
Bonding and Structure 12
Electron Electron Distribution in Distribution in
MoleculesMolecules• Electron distribution
is depicted with Lewis electron dot structures
• Electrons are distributed as:
• shared or BOND PAIRS and
• unshared or LONE PAIRS.
G. N. Lewis 1875 - 1946
Bonding and Structure 13
Bond and Lone Pairs• Electrons are distributed as shared or BOND
PAIRS and unshared or LONE PAIRS.
••H Cl
••••
This is a LEWIS ELECTRON DOT structure.
shared or bond pair
Unshared orlone pair (LP)
Bonding and Structure 14
• This observation is called the OCTET RULE
Rules of Lewis StructuresRules of Lewis Structures• No. of valence electrons of an atom =
Group number
• Except for H (and atoms of 3rd and higher periods),
#Bond Pairs + #Lone Pairs = 4
• For Groups 5A-7A (N - F), no. of BOND PAIRS = 8 - group No.
• For Groups 1A-4A (Li - C), no. of BOND PAIRS = group number
Bonding and Structure 15
2. Count valence electrons H = 1 and N = 5 Total = (3 x 1) + 5 = 8 electrons or
1. Decide on the central atom; never H. Central atom is atom of lowest affinity for electrons. In ammonia, N is central
Building a Dot StructureAmmonia, NH3
4 pairs
Bonding and Structure 16
4. Remaining electrons form LONE PAIRS to complete octet as needed.
3. Form a sigma bond between the central atom and surrounding atoms.
H HHN
Building a Dot Structure
H•• HHN
3 BOND PAIRS and 1 LONE PAIR.
Note that N has a share in 4 pairs (8 electrons), while each H shares 1 pair.
Bonding and Structure 17
Step 2. Count valence electrons S = 6 3 x O = 3 x 6 = 18 Negative charge = 2
TOTAL = 6 + 18 + 2 = 26 e- or 13 pairs
Step 1. Central atom = S
10 pairs of electrons are left.
Sulfite ion, SO32-
Step 3. Form sigma bonds
O O
O
S
Bonding and Structure 18
Thanks to all
Bonding and Structure 19
Remaining pairs become lone pairs, first on outside atoms
then on central atom.
Sulfite ion, SO32- (2)
Each atom is surrounded by an octet of electrons.
••O O
O
S
••
••
•• ••
••
••
••••
••
NOTE - must add formal charges (O-, S+) for complete dot diagram
Bonding and Structure 20
Carbon Dioxide, CO2
1. Central atom = __C____2. Valence electrons = _16_ or _8_ pairs3. Form sigma bonds.
O OC
••O OC
•• ••
••••••
This leaves __6__ pairs.4. Place lone pairs on outer atoms.
Bonding and Structure 21
••O OC
•• ••
••••••
Carbon Dioxide, CO2 (2)4. Place lone pairs on outer atoms.
••O OC
•• ••
••••••
••O OC
•• ••
••
The second bonding pair forms a pi () bond.
5. To give C an octet, form DOUBLE BONDS between C and O.
Bonding and Structure 22
SO3
H2CODouble and even triple bonds are commonly observed for C, N, P, O, and S
••O OC
•• ••
••C2F4
Bonding and Structure 23
Sulfur Dioxide, SO21. Central atom = S2. Valence electrons = 6 + 2*6 = 18 electrons
or 9 pairs••O OS
••
••
••
••••••
••O OS
••
••
••
••••••
bring inleft pair
OR bring inright pair
3. Form pi () bond so that S has an octet — note that there are two ways of doing this.
Bonding and Structure 24
Sulfur Dioxide, SO2
••O OS
••
••
••
••••••
bring inleft pair
OR bring inright pair
••O OS
••
••
••
••••
••O OS••
••
••
••••
Equivalent structurescalled:
RESONANCE STRUCTURES
The proper Lewis structure is a HYBRID of the two.
A BETTER representation of SO2 is made by forming 2 double bonds
O = S = OEach atom has - OCTET - formal charge = 0
Bonding and Structure 25
Urea (NHUrea (NH22))22COCO1. Number of valence electrons = 24 e-2. Draw sigma bonds.
CN N HHH
H
O
4. Place remaining electron pairs on oxygen3. Complete C atom octet with double bond.
CN N HHH
H
O
Leaves 24 - 14 = 10 e- pairs.
and nitrogen atoms.
Bonding and Structure 26
Violations of the Octet RuleUsually occurs with:
Boron
BF3 SF4
elements of higher periods.
Bonding and Structure 27
Boron Trifluoride• Central atom = B• Valence electrons = 3 + 3*7 = 24 or electron pairs = 12• Assemble dot structure
F••
••
••
F
FB
••
••
••
••
••
••
The B atom has a share in only 6 electrons (or 3 pairs). B atom in many molecules is electron deficient.
Bonding and Structure 28
Sulfur Tetrafluoride, SF4
• Central atom = S• Valence electrons = 6 + 4*7 = 34 e-
or 17 pairs.• Form sigma bonds and distribute
electron pairs.
F
••
••
••
FF
S••
••••
••
•• F
••
••
••
••
•• 5 pairs around the S 5 pairs around the S atom. A common atom. A common occurrence outside the occurrence outside the 2nd period. 2nd period.
Bonding and Structure 29
Explanation of the failure of octet rule
Sidgwick’s rule of maximum covalency
• It is not essential to have 8 electron surrounding in an atom
• Maximum covalency depends on its period
period Maximum shared electron
Atom
n=1 2 H
n=2 4 Li, F
n=3 6 Na, Mg
n=4 8 k
Bonding and Structure 30
Sugden’s view of singlet linkages
Octet rule never violated
In PCl5, 3 Cl atom linked by normal covalent bond and others 2 by sharing only 1 electron called singlet linkage
Formal Charges• Formal charge is the charge calculated
for an atom in a Lewis structure on the basis of an equal sharing of bonded electron pairs.
Nitric acid
.... ::
....HH OO
OO
OO
NN
::
::....
....
• We will calculate the formal charge for each atom in this Lewis structure.
Formal charge of HFormal charge of H
Nitric acid
.... ::
....HH OO
OO
OO
NN
::
::....
....
• Hydrogen shares 2 electrons with oxygen.• Assign 1 electron to H and 1 to O.• A neutral hydrogen atom has 1 electron.• Therefore, the formal charge of H in nitric
acid is 0.
Formal charge of HFormal charge of H
Nitric acid
.... ::
....HH OO
OO
OO
NN
::
::....
....
• Oxygen has 4 electrons in covalent bonds.• Assign 2 of these 4 electrons to O.• Oxygen has 2 unshared pairs. Assign all 4 of
these electrons to O.• Therefore, the total number of electrons assigned
to O is 2 + 4 = 6.
Formal charge of OFormal charge of O
Nitric acid
.... ::
....HH OO
OO
OO
NN
::
::....
....
• Electron count of O is 6.• A neutral oxygen has 6 electrons.• Therefore, the formal charge of O is 0.
Formal charge of OFormal charge of O
Nitric acid
.... ::
....HH OO
OO
OO
NN
::
::....
....
• Electron count of O is 6 (4 electrons from unshared pairs + half of 4 bonded electrons).
• A neutral oxygen has 6 electrons.• Therefore, the formal charge of O is 0.
Formal charge of OFormal charge of O
Nitric acid
.... ::
....HH OO
OO
OO
NN
::
::....
....
• Electron count of O is 7 (6 electrons from unshared pairs + half of 2 bonded electrons).
• A neutral oxygen has 6 electrons.• Therefore, the formal charge of O is -1.
Formal charge of OFormal charge of O
Nitric acid
.... ::
....HH OO
OO
OO
NN
::
::....
....
• Electron count of N is 4 (half of 8 electrons in covalent bonds).
• A neutral nitrogen has 5 electrons.• Therefore, the formal charge of N is +1.
Formal charge of NFormal charge of N
––
Nitric acid
.... ::
....HH OO
OO
OO
NN
::
::....
....
• A Lewis structure is not complete unless formal charges (if any) are shown.
Formal chargesFormal charges
––++
Formal ChargeFormal Charge
Formal charge = Formal charge =
group numbergroup numberin periodic tablein periodic table
number ofnumber ofbondsbonds
number ofnumber ofunshared electronsunshared electrons–– ––
An arithmetic formula for calculating formal charge.An arithmetic formula for calculating formal charge.
"Electron counts""Electron counts" and formal and formal charges in NHcharges in NH44
+ + and BFand BF44--
11
44
NN
HH
HH HH
HH
++77
44
....
BBFF
FF
FF
FF
....
............:: ::
:: ::
:: ::
....
––
Resonance
Resonance
In chemistry, resonance or mesomerism is a way of
describing delocalized electrons within certain molecules
or polyatomic ions where the bonding cannot be
expressed by one single Lewis formula. A molecule or ion
with such delocalized electrons is represented by several
contributing structures (also called resonance structures
or canonical forms).
Bonding and Structure 44
General characteristics of resonanceMolecules and ions with resonance (also called
mesomerism) have the following basic characteristics:•They can be represented by several correct Lewis formulas, called
"contributing structures", "resonance structures“ or "canonical forms".
However, the real structure is not a rapid interconversion of contributing
structures. Several Lewis structures are used together, because none of
them exactly represents the actual structure. To represent the
intermediate, a resonance hybrid is used instead.
•The contributing structures are not isomers. They differ only in the
position of electrons, not in the position of nuclei.
Bonding and Structure 45
• Each Lewis formula must have the same number of valence electrons (and thus the same total charge), and the same number of unpaired electrons, if any.
• Bonds that have different bond orders in different
contributing structures do not have typical bond lengths. Measurements reveal intermediate bond lengths.
• The real structure has a lower total potential energy than each of the contributing structures would have. This means that it is more stable than each separate contributing structure would be.
General characteristics of resonance
Bonding and Structure 46
Contributing structures of some ion
•same atomic positions
•differ in electron positions
more stable more stable Lewis Lewis
structurestructure
less stable less stable Lewis Lewis
structurestructure
........
CC OO NN OOHH
HH
HH.... ::....
++ ––....
....CC OO NN OOHH
HH
HH
....::....
Resonance Structures of Methyl Nitrite
•same atomic positions
•differ in electron positions
more stable more stable Lewis Lewis
structurestructure
less stable less stable Lewis Lewis
structurestructure
........
CC OO NN OOHH
HH
HH.... ::....
++ ––....
....CC OO NN OOHH
HH
HH
....::....
Resonance Structures of Methyl Nitrite
• Electrons in molecules are often delocalizedbetween two or more atoms.
• Electrons in a single Lewis structure are assigned to specific atoms-a single Lewis structure is insufficient to show electron delocalization.
• Composite of resonance forms more accurately depicts electron distribution.
Why Write Resonance Structures?
•Ozone (O3)
–Lewis structure of
ozone shows one
double bond and
one single bond
Expect: one short bond and one Expect: one short bond and one long bondlong bond
Reality: bonds are of equal length Reality: bonds are of equal length (128 pm)(128 pm)
Example
OO OO••••OO••••••••••••
••••••••––++
•Ozone (O3)
–Lewis structure of
ozone shows one
double bond and
one single bond
Resonance:Resonance:
Example
OO OO••••OO••••••••••••
••••••••––++
OO OO••••OO••••••••••••
••••••••––++
OO OOOO••••••••••••
••••••••–– ++
••••
Bonding and Structure 52
Bonding and Structure 53
Formal charge = Group no. - 1/2 (no. bond electrons)
- (no. of LP electrons)
Formal Atom Charges• Atoms in molecules often bear a charge (+ or -).
• The most important dominant resonance structure of a molecule is the one with formal charges as close to 0 as possible.
Bonding and Structure 54
04 - (1/2)(8) - 0 =
6 - (1/2)(4) - 4 = 0
Carbon Dioxide, CO2
At OXYGEN
O C O••
• •••
• •
At CARBON
Bonding and Structure 55
C atom charge is 0.
6 - (1/2)(6) - 2 = +1
6 - (1/2)(2) - 6 = -1
Carbon Dioxide, CO2 (2)
O C O••
• •••
• •
An alternate Lewis structure is:
AND the corresponding resonance form
+
O C O••
• •••
• •
+
Bonding and Structure 56
• REALITY: Partial charges calculated by CAChe molecular modeling system (on CD-ROM).
+1.46-0.73 -0.73
Carbon Dioxide, CO2 (3)Which is the predominant resonance structure?
O C O••
• •••
• •
ORO C O•
•• •
••
• •
O C O••
• •••
• •
+
+
Answer ?Form without formal charges isBETTER - no +ve charge on O
Bonding and Structure 57
Boron Trifluoride, BF3
F••
••
••
F
FB
••
••
••
••
••
••
What if we form a B—F double bond to satisfy the B atom octet?
Bonding and Structure 58
Boron Trifluoride, BF3 (2)
• To have +1 charge on F, with its very high electron affinity is not good. -ve charges best placed on atoms with high EA.
• Similarly -1 charge on B is bad• NOT important Lewis structure
fc = 7 - 2 - 4 = +1 Fluorine
F••
••
F
FB
••
••
••
••
••
••
fc = 3 - 4 - 0 = -1 Boron
+
Bonding and Structure 59
A. S=C=N
Thiocyanate ion, (SCN)-Which of three possible resonance structuresis most important?
-0.52 -0.32-0.16
Calculated partial chargesANSWER:C > A > B
C. S-C N
B. S=C - N