Chemical Bonding
and
Molecular Structure
(Chapter 9)
• Ionic vs. covalent bonding
• Molecular orbitals and the covalent bond (Ch. 10)
• Valence electron Lewis dot structures
octet vs. non-octet
resonance structures
formal charges
• VSEPR - predicting shapes of molecules
• Bond properties
electronegativity
polarity, bond order, bond strength
20 Oct 97
Bonding and structure (2)
1
Rules for making Lewis dot structures
1. Count no. of valence electrons
(- don’t forget to include the charge on molecular ions!)
— 2 for # of PAIRS
2. Place a bond pair (BP) between connected atoms
3. Complete octets by using rest of e- as lone pairs (LP)
4. For atoms with <8 e-, make multiple bonds to complete octets
5. Assign formal charges : fc = Z - (#BP/2) - (#LP)
Indicate equivalent (RESONANCE) structures
6. Structures with smaller formal charges are preferred
- consider non-octet alternatives (esp. for 3rd, 4th row)
•OCTET RULE: #Bond Pairs + #Lone Pairs = 4
(except for H and atoms of 3rd and higher periods)
#lone pairs at central atom in AXn = {(#e-) - 8*n}/2
20 Oct 97
Bonding and structure (2)
2
Sulfur Dioxide, SO2
Rules 1-3 
O—S —O
OR bring in
right pair
bring in
left pair
•
•
••
••
••
O
S
O
••
••
•
•
O
••
••
••
•
•
O
S
+
—
••
•
•
••
•
•
••
••
O— +S
•
•
O
These equivalent structures
are called:
••
The proper Lewis structure
is a HYBRID of the two.
RESONANCE
STRUCTURES.
Each atom has OCTET . .
. . . BUT there is a +1 and -1 formal charge
20 Oct 97
Bonding and structure (2)
3
SO2 (2)
Alternate Lewis structure for SO2 uses 2 double bonds
O=S=O
Sulfur does not obey OCTET rule
BUT the formal charge = 0
This is better structure than O=S+-Osince it reduces formal charge (rule 6).
3rd row S atom can have 5 or 6 electron pairs
NB: # of central atom lone pairs = (3*6 -8*2)/2 = 1
in both O=S+-O- and O=S=O structures
20 Oct 97
Bonding and structure (2)
4
Thiocyanate ion, (SCN)
-
Which of three possible resonance structures
is most important?
A.
S=C=N
B.
S=C - N
C.
S-C N
Calculated partial charges
20-0.52
Oct 97
-0.16
ANSWER:
C>A>B
-0.32
Bonding and structure (2)
5
MOLECULAR GEOMETRY
VSEPR
• Valence Shell Electron Pair
Repulsion theory.
• Most important factor in
determining geometry is relative
repulsion between electron pairs.
20 Oct 97
Bonding and structure (2)
Molecule
adopts the
shape that
minimizes the
electron pair
repulsions.
6_VSEPR.mov
6
No. of e- Pairs
Around Central
Atom
2
Example
F—Be—F
Geometry
CAChe
image
linear
180o
F
3
F
B
F
planar trigonal
120o
H
4
C
H
20 Oct 97
109o
tetrahedral
H
H
Bonding and structure (2)
7
Structure Determination by VSEPR
Ammonia, NH3
There are 4 electron pairs at the corners of a
tetrahedron.
lone pair of electrons
in tetrahedral position
••
H
N
H
H
N
H
H
H
The ELECTRON PAIR GEOMETRY is tetrahedral.
20 Oct 97
Bonding and structure (2)
8
VSEPR - ammonia
lone pair of electrons
in tetrahedral position
Ammonia, NH3
N
H
H
H
Although the electron pair geometry is tetrahedral . . .
. . . the MOLECULAR GEOMETRY
— the positions of the atoms
— is PYRAMIDAL.
20 Oct 97
Bonding and structure (2)
9
AXnEm notation
• a good way to distinguish between
electron pair and molecular geometries
is the AXnEm notation
where:
A - atom whose local geometry is of interest
(typically the CENTRAL ATOM)
Xn - n atoms bonded to A
Em - m lone pair electrons at A
NH3 is AX3E system  pyramidal
(NB this notation not used by Kotz)
20 Oct 97
Bonding and structure (2)
10
VSEPR - water
Water, H2O
••
1. Draw electron dot structure
H-O-H
••
2. Count BP’s and LP’s = 4
3. The 4 electron pairs are at the
corners of a tetrahedron.
O
H
H
20 Oct 97
The electron pair
geometry is
TETRAHEDRAL.
Bonding and structure (2)
11
VSEPR - water (2)
••
H-O-H
••
O
Although the electron
pair geometry is
H TETRAHEDRAL . . .
H
. . . the molecular
geometry is bent.
H2O - AX2E2 system - angular geometry
20 Oct 97
Bonding and structure (2)
12
VSEPR - formaldehyde
Formaldehyde, CH2O
• •
1. Draw electron dot structure
•O•
2. Count BP’s and LP’s:
H C H
At Carbon there are 4 BP but . . .
3. These are distributed in ONLY 3 regions.
Double bond electron pairs are in same region.
There are 3 regions of electron density
Electron repulsion places them at the corners of a
planar triangle.
• •
Both the electron pair geometry and the
• O•
molecular geometry are PLANAR TRIGONAL
o bond angles.

120
C
H
H
20 Oct 97
H2CO at the C atom is an AX3 species
Bonding and structure (2)
13
VSEPR - Bond Angles
Methanol, CH3OH
H
••
H—C—O—H
••
H
Angle 1
Angle 2
Define bond angles 1 and 2
Angle 1 = H-C-H = ?
Angle 2 = H-O-C = ?
Answer: 109o because both the C and O
atoms are surrounded by 4 electron pairs.
6_CH3OH.mov
AXnEm designation ?
at C
at O
20 Oct 97
Bonding and structure (2)
AX4 = tetrahedral
AX2E2 = bent
14
VSEPR - bond angles (2)
Acetonitrile, CH3CN
H
Define bond angles 1 and 2
N
H—C—C
1
Angle 2 = 180
? o
H
••
109o
Angle 1 = ?
2
Why ? :
The CH3 carbon is surrounded by 4 bond charges
The CN carbon is surrounded by 2 bond charges
AXnEm designation ?
at CH3 carbon AX4 = tetrahedral
at CN carbon
20 Oct 97
AX2 = linear
Bonding and structure (2)
15
What about:
STRUCTURES WITH
CENTRAL ATOMS
THAT DO NOT OBEY
THE OCTET RULE ?
PF5
BF3
SF4
20 Oct 97
Bonding and structure (2)
16
Geometry for non-octet species
also obey VSEPR rules
Consider boron trifluoride, BF3
The B atom is surrounded by only 3
electron pairs.
Bond angles are 120o
••
•
•
•
•
F
••
•
•
F
••
B
•
•
•
•
F
••
Molecular Geometry is
planar trigonal
BF3 is an AX3 species
20 Oct 97
Bonding and structure (2)
17
Compounds with 5 or 6 Pairs
Around the Central Atom
90
F
F
P
Trigonal bipyramid
F
120 5 electron pairs
F
AX5 system
F
90
6 electron pairs
F
F
S
F
F
20 Oct 97
6_VSEPR.mov
Bonding and structure (2)
Octahedron
F
F
90
AX6 system
18
Sulfur Tetrafluoride, SF4
••
••
•• F
••
•••
••
••F S F•
••
••
•• F••
••
Number of valence e- = 34
No. of S lone pairs =
{17 - 4 b.p. - 3x4 l.p.(F)}
= 1 lone pair on S
There are 5 (BP + LP)
e- pairs around the S
THEREFORE:
electron pair geometry ?
= trigonal bipyramid
F
•
•
S
F
OR
F
F
••
F
F
F
S
F
AX4E system. Molecular geometry ?
20 Oct 97
Bonding and structure (2)
19
Sulfur Tetrafluoride, SF4 (2)
90
axial
equatorial
••
F
S
F
120
F
F
Lone pair is in the equatorial position because it
requires more room than a bond pair.
Molecular geometry of SF4 is “see-saw”
Q: What is molecular geometry of SO2 ?
20 Oct 97
Bonding and structure (2)
20
Bonding with Hybrid Atomic Orbitals
- Carbon prefers to make 4 bonds as in CH4
But atomic carbon has an s2p2 configuration
Why can it make more than 2 bonds ?
6_CH4.mov
4 C atom orbitals hybridize to form four
equivalent sp3 hybrid atomic orbitals.
20 Oct 97
Bonding and structure (2)
21
Orbital Hybridization
BONDS
SHAPE
HYBRID REMAIN
e.g.
s2p2 
2
linear
{2 x sp &2 p’s}
C2H2
3
trigonal
planar
{3 x sp2 & 1 p}
C2H4
4
tetrahedral {4 xsp3 }
20 Oct 97
Bonding and structure (2)
CH4
22
Multiple Bonds
s and p Bonding in C2H4

2s
2p
C atom orbitals are COMBINED
(= re-hybridized) to form orbitals
better suited for BONDING
•
• The 3 sp2 hybrid orbitals
are used to make the C-C
and two C-H s bonds
6_C2H4-sg.mov
6_C2H4-pi.mov
120
The extra p orbital
electron on each C atom
overlaps the p orbital on
the neighboring atom to
form the p bond.
H
H
6_C2H4.mov
20 Oct 97
p
orbital
3 sp2
hybrid
orbitals
Bonding and structure (2)
C
C
H
sp2
H
23
Consequences of Multiple Bonding
Restricted rotation around C=C bond in
1-butene = CH2=CH-CH2-CH3.
E (kJ/mol)
See Butene.Map in ENER_MAP in CAChe models.
233
27
-180
20 Oct 97
0
180
C-C=C angle (o)
P. 475 - Photo-rotation
about double bonds
lets us see !!
Bonding and structure (2)
24
Bond Properties
• What is the effect of bonding and structure on
molecular properties ?
- bond order
- bond length
- bond strength
- bond polarity
- MOLECULAR polarity
Buckyball in HIV-protease, see page 107
20 Oct 97
Bonding and structure (2)
25
Bond Order
• the number of bonds between a pair of atoms.
H
single
BO = 1
1 s
20 Oct 97
H
H
C
C
triple, BO = 3
1 s and 2 p
C
CH2CHCN
Acrylonitrile
N
double, BO = 2
1 s and 1 p
Bonding and structure (2)
26
Bond Order (2)
Fractional bond orders occur in molecules with
resonance structures.
Consider NO2-
••
O
••
••
N
•• •
O•
••
••••
O
••
••
N
••
O
••
Total # of e - pairs used for a type of bond
Bond order =
Total # of bonds of that type
3
(e
pairs
in
N-O
bonds)
Bond order in NO2 =
2 (N - O bonds)
N-O bond order in NO2- = 1.5
-
20 Oct 97
Bonding and structure (2)
27
Bond Order and Bond Length
Bond order is related to two important bond properties:
(a) bond strength
as given by DE
110 pm
745 kJ
(b) Bond length
- the distance between
the nuclei of two bonded
atoms.
20 Oct 97
Bonding and structure (2)
414 kJ 123 pm
Formaldehye
28
Bond Length
- depends on size of bonded atoms:
Molecule
H- F
H- Cl
H- I
- depends on bond order.
Molecule
CH3C- OH
O=C=O
C O
20 Oct 97
R(H-X)
104 pm
131 pm
165 pm
R(C-O)
141 pm
132 pm
119 pm
Bonding and structure (2)
29
Bond Strength
• Bond Dissociation energy (DE) - energy required to
break a bond in gas phase.
• See Table 9.5
BOND
STRENGTH (kJ/mol) LENGTH (pm)
H—H
436
74
C—C
C=C
CC
347
611
837
154
134
121
NN
946
110
The GREATER the number of bonds (bond order)
the HIGHER the bond strength and the SHORTER the bond.
20 Oct 97
Bonding and structure (2)
30
Bond Strength (2)
Bond
Order
Length
Strength
HO—OH
1
149 pm
210 kJ/mol
O=O
2
121
498 kJ/mol
128
? kJ/mol
303
••
O
••
••
O
•• •
O•
••
1.5
HOW TO CALCULATE ?
O3 (g)  3 O(g)
Hrxn = {3xHf(O) - Hf(O3)} = {3x249.2 - 142.7} = 605 kJ/mol
2 O-O bonds in O3  DE (O3) = 605/2 = 302.5 kJ/mol
20 Oct 97
Bonding and structure (2)
31
Bond Polarity
+
H
-
••
Cl ••
••
HCl is POLAR because it has a positive
end and a negative end (partly ionic).
Polarity arises because Cl has a greater
share of the bonding electrons than H.
Calculated charge by CAChe:
H (red)
is +ve (+0.20 e-)
Cl (yellow) is -ve (-0.20 e-).
(See PARTCHRG folder in MODELS.)
20 Oct 97
Bonding and structure (2)
32
+
Bond Polarity (2)
H
-
••
Cl ••
••
• Due to the bond polarity, the H—Cl bond
energy is GREATER than expected for a
“pure” covalent bond.
BOND
ENERGY
“pure” bond
339 kJ/mol calculated
real bond
432 kJ/mol measured
Difference
92 kJ/mol.
This difference is the contribution of IONIC bonding
It is proportional to the difference in
ELECTRONEGATIVITY, c.
20 Oct 97
Bonding and structure (2)
33
Electronegativity, c
c is a measure of the ability of an atom in a
molecule to attract electrons to itself.
Concept proposed by
Linus Pauling (1901-94)
Nobel prizes:
Chemistry (54), Peace (63)
See p. 425; 008vd3.mov (CD)
20 Oct 97
Bonding and structure (2)
34
4
N
3.5
Cl
C
3
2.5
F
O
H
Si
2
P
S
Electronegativity, c
Figure 9.7
1.5
1
0.5
0
1
2
3
4
5
6
7
8
9
10
11
12
13
14 15
16
17
18
• F has maximum c.
• Atom with lowest c is the center atom in most
molecules.
• Relative values of c determines BOND
POLARITY (and point of attack on a molecule).
20 Oct 97
Bonding and structure (2)
35
Bond Polarity
Which bond is more polar ? (has larger bond DIPOLE)
O—H
O—F
c H
c(A) - c(B) 3.5 - 2.1
3.5 - 4.0
0.5
c
1.4
c(O-H) > c(O-F)
Therefore OH is more polar than OF
2.1
O F
3.5 4.0
Also note that polarity is “reversed.”
O—H
- +
20 Oct 97
O—F
+ -
Bonding and structure (2)
36
Molecular Polarity
• Molecules—such as HCl and H2O—
can be POLAR (or dipolar).
• They have a DIPOLE MOMENT.
• Polar molecules turn to align their
dipole with an electric field.
POSITIVE
20 Oct 97
H—Cl


Bonding and structure (2)
NEGATIVE
37
Predicting molecular polarity
A molecule will be polar ONLY if
a) it contains polar bonds
AND
b) the molecule is NOT “symmetric”
Symmetric molecules
20 Oct 97
Bonding and structure (2)
38
Molecular Polarity: H2O
••
O
••
H
H
polar
O
H
H
+
Water is polar because:
a) O-H bond is polar
b) water is non-symmetric
The dipole associated with polar H2O
is the basis for absorption of microwaves
used in cooking with a microwave oven
20 Oct 97
Bonding and structure (2)
39
Carbon Dioxide
-0.73 +1.46 -0.73
• CO2 is NOT polar even
though the CO bonds
are polar.
• Because CO2 is
symmetrical the BOND
polarity cancels
The positive C atom is why
water attaches to CO2
CO2 + H2O  H2CO3
20 Oct 97
Bonding and structure (2)
40
Molecular Polarity in
NON-symmetric molecules
F
B +ve
F -ve
B
F
B
F
F
B—F bonds are polar
molecule is symmetric
BF3 is NOT polar
20 Oct 97
Atom Chg. c
H
F
B
H
F
+ve
+ve
-ve
2.0
2.1
4.0
B—F, B—H bonds polar
molecule is NOT symmetric
HBF2 is polar
Bonding and structure (2)
41
Fluorine-substituted Ethylene: C2H2F2
C—F bonds are MUCH more polar than C—H bonds.
c(C-F) = 1.5, c(C-H) = 0.4
CIS isomer
• both C—F bonds on same side
 molecule is POLAR.
TRANS isomer
• both C—F bonds on opposite side
 molecule is NOT POLAR.
20 Oct 97
Bonding and structure (2)
42
Chemical Bonding
and
Molecular Structure
(Chapter 9)
• Ionic vs. covalent bonding
• Molecular orbitals and the covalent bond (Ch. 10)
• Valence electron Lewis dot structures
octet vs. non-octet
resonance structures
formal charges
• VSEPR - predicting shapes of molecules
• Bond properties
electronegativity
polarity, bond order, bond strength
20 Oct 97
Bonding and structure (2)
43