Ch. 6 – Molecular Structure
I. Lewis Diagrams
(p. 184-189)
I
II
III
A. Octet Rule

Remember…
 Most atoms form bonds in order to
have 8 valence electrons.
A. Octet Rule

F
F
 Hydrogen  2 valence e
F
B
F
 Groups F
1,2,3 get
2,4,6
valence e
S
F
H
N
O
O
H
 Expanded octet

more
than
8
F
Very
unstable!!
valence
e
(e.g.
S,
P,
Xe)
F
F
Exceptions:
-
-
-
 Radicals  odd # of valence e-
B. Drawing Lewis Diagrams

Find total # of valence e-.

Arrange atoms - singular atom is
usually in the middle.

Form bonds between atoms (2 e-).

Distribute remaining e- to give each
atom an octet (recall exceptions).

If there aren’t enough e- to go around,
form double or triple bonds.
B. Drawing Lewis Diagrams
 CF4
1 C × 4e- = 4e4 F × 7e- = 28e32e- 8e24e-
F
F C F
F
B. Drawing Lewis Diagrams
 BeCl2
1 Be × 2e- = 2e2 Cl × 7e- = 14e16e- 4e12e-
Cl Be Cl
B. Drawing Lewis Diagrams
 CO2
1 C × 4e- = 4e2 O × 6e- = 12e16e- 4e12e-
O C O
C. Polyatomic Ions

To find total # of valence e-:
 Add 1e- for each negative charge.
 Subtract 1e- for each positive
charge.

Place brackets around the ion and
label the charge.
C. Polyatomic Ions
 ClO4-
1 Cl × 7e- = 7e4 O × 6e- = 24e31e+ 1e32e- 8e24e-
O
O Cl O
O
C. Polyatomic Ions
 NH4+
1 N × 5e- = 5e4 H × 1e- = 4e9e- 1e8e- 8e0e-
H
H N H
H
D. Resonance Structures
Molecules that can’t be correctly
represented by a single Lewis
diagram.
 Actual structure is an average of all
the possibilities.
 Show possible structures separated
by a double-headed arrow.

D. Resonance Structures
 SO3
O
O S O
O
O S O
O
O S O
Ch. 6 – Molecular Structure
II. Molecular
Geometry
(p. 197-200)
I
II
III
A. VSEPR Theory

Valence Shell Electron Pair
Repulsion Theory

Electron pairs orient themselves in
order to minimize repulsive forces.
A. VSEPR Theory

Types of e- Pairs
 Bonding pairs - form bonds
 Lone pairs - nonbonding e-
Lone pairs repel
more strongly than
bonding pairs!!!
A. VSEPR Theory

Lone pairs reduce the bond angle
between atoms.
Bond Angle
B. Determining Molecular Shape

Draw the Lewis Diagram.

Tally up e- pairs on central atom.
 double/triple bonds = ONE pair

Shape is determined by the # of
bonding pairs and lone pairs.
Know the 8 common shapes
& their bond angles!
C. Common Molecular Shapes
2 total
2 bond
0 lone
BeH2
LINEAR
180°
C. Common Molecular Shapes
3 total
3 bond
0 lone
BF3
TRIGONAL PLANAR
120°
C. Common Molecular Shapes
3 total
2 bond
1 lone
SO2
BENT
<120°
C. Common Molecular Shapes
4 total
4 bond
0 lone
CH4
TETRAHEDRAL
109.5°
C. Common Molecular Shapes
4 total
3 bond
1 lone
NH3
TRIGONAL PYRAMIDAL
107°
C. Common Molecular Shapes
4 total
2 bond
2 lone
H2O
BENT
104.5°
D. Examples

PF3
4 total
3 bond
1 lone
F P F
F
TRIGONAL
PYRAMIDAL
107°
D. Examples

CO2
2 total
2 bond
0 lone
O C O
LINEAR
180°
Ch. 6 – Molecular Structure
III. Polarity &
IMF
(p. 204-207)
I
II
III
A. Dipole Moment

Direction of the polar bond in a
molecule.

Arrow points toward the more e-neg
atom.
+

H
Cl

B. Determining Molecular Polarity

Depends on:
 dipole moments
 molecular shape
B. Determining Molecular Polarity

Nonpolar Molecules
 Dipole moments are symmetrical
and cancel out.
F
BF3
B
F
F
B. Determining Molecular Polarity

Polar Molecules
 Dipole moments are asymmetrical
and don’t cancel .
O
H2O
H
H
net
dipole
moment
B. Determining Molecular Polarity

Therefore, polar molecules have...
 asymmetrical shape (lone pairs) or
 asymmetrical atoms
H
CHCl3
Cl
Cl
Cl
net
dipole
moment
Dipole-Dipole Forces

Attractive forces between polar
covalent molecules
London (Dispersion) Forces
Attractive forces between the electron
clouds of large molecules in large
quantity
 Larger mass = Larger London Forces

Hydrogen Bonding

Special dipole-dipole attraction that
involves H bonded with high
electronegative elements N, O, or F