Shape and Bonding
1.  Determine Lewis structure.
2.  Use VSEPR to determine electron pair
geometry.
3.  Determine molecular geometry.
4. Is the molecule polar?
Dipole moment
µ = Q•r
To answer this, use knowledge of:
  Molecular shape
  Bond Polarity
Determine Molecular Polarity
Mary J. Bojan
Chem 110
1
Sharing of electrons in bonds is not always equal
Covalent bond
identical atoms:
Polar bond
different atoms:
•  large difference:
ionic
N2
Cl2
H2
+⎯→
H⎯Cl
δ+ δ−
Bond type is determined by differences in electronegativity (ΔEN).
Na+
0.93
Cl−
3.2
•  intermediate values: polar covalent
H — Cl
2.2 3.2
•  no difference:
covalent
Cl — Cl
approximate cutoff:
ΔEN ≥ 2.0
ΔEN < 2.0
ionic
covalent
Mary J. Bojan
Chem 110
ΔEN =
ΔEN =
2
BOND POLARITY
AND DIPOLE MOMENTS
One measure of real charges in molecules is dipole moment (µ)
For a diatomic molecule: µ = Q r
| ⎯ r ⎯ |
•⎯⎯⎯•
+Q
−Q
unit = debye (D) =
= 3.33 x10−30 C-m
For two charges (+1 and −1) separated by 1 Å
µ = 4.79 D
Units of Q = 1.6 x10−19 C = charge of an electron
Mary J. Bojan
Chem 110
3
DIPOLE TRENDS in HX
HX
bond
length Å
µ(exp)
D
µ(ionic)
D
% ionic
HF
0.92
1.82
4.41
41%
HCl
1.27
1.08
6.08
18%
HBr
1.41
0.82
6.75
12%
HI
1.61
0.44
7.71
6%
polar covalent
H⎯X
δ+ δ−
Consider HCl:
the bond length is 1.27 Å
if there were full charge separation:
µ ionic = 1.27 Å x 4.79 D /Å = 6.08 D
HCl molecule: dipole moment = 1.08 D
This means the charge separation is only
(1.08/6.08) = 0.18 of the full charge
Mary J. Bojan
Chem 110
Sample calculation:
µionic = 4.79D/ Å *(bond length)
For HCl, = 4.79D/ Å * 1.27 Å = 6.08 D
% ionic = µ(exp)/ µ(ionic) x 100%
= 1.08D / 6.08D x 100% = 18%
4
DIPOLE TRENDS in HX
HX
bond
length Å
µ(exp)
D
ΔEN
HF
0.92
1.82
1.9
HCl
1.27
1.08
0.9
HBr
1.41
0.82
0.7
HI
1.61
0.44
0.4
Trends: µexp increases
−
−
For a linear diatomic molecule
µ depends on:
For a polyatomic molecule
µ depends on:
bond polarities
molecular geometry
Mary J. Bojan
Chem 110
5
Oxidation Numbers, Formal Charges and
Partial Charges
Consider HCl:
+1 −1
H-Cl
0
0
H-Cl
oxidation formal
numbers charges
+δ -δ
HCl δ = 0.18
partial charges
(real)
Partial charges: "real" charges on atoms in molecules.
Explicitly include electronegativity differences.
dipole moment (µ) is one measure of real charges
Mary J. Bojan
Chem 110
6
Polarity of Molecules
A molecule is polar if there is a NET charge separation between
two "ends" of the molecule: molecule has a negative "end" and a
positive "end".
To have a net dipole:
1. polar bonds
2. geometry where bond dipoles DO NOT cancel
________________________________
To determine the polarity of a molecule that has more
than 2 atoms:
1.  find molecular shape (3D)
2.  find "bond" dipoles
3.  use vector "analysis" to find net molecular dipole
Mary J. Bojan
Chem 110
7
Example: CO2
Draw the Lewis structure:
O=C=O
EPG:
MG:
bond dipoles?
electronegativity:
O is more electronegative than C
Net dipole moment?
Mary J. Bojan
Chem 110
8
Example: H2O
Draw the Lewis Structure
EPG:
MG:
Bond Dipoles?
electronegativity:
O is more electronegative than H
Net dipole moment?
Mary J. Bojan
Chem 110
9
Example: NH3
EPG: tetrahedral
MG : trigonal pyramid
Bond polarity
electronegativity:
N>H
H
N
H
H
Net dipole moment?
Mary J. Bojan
Chem 110
10
More Examples
CF4
F
EPG:
MG:
Bond dipoles?
F
C
F
F
Net dipole moment?
electronegativity: F > C
CH3Cl
EPG:
MG:
electronegativity: H ≈ C
Cl > C
Mary J. Bojan
Chem 110
Net dipole
moment?
11
Molecular Structure
1. Determine Lewis structure.
2. Use VSEPR to determine electron pair geometry.
3. Determine molecular geometry.
4. Determine bond polarities: is the molecule polar?
CO2 is not polar
Is SO2 polar?
CF4 is not polar
Is SF4 polar?
Mary J. Bojan
Chem 110
12