Chapter 9
Molecular Geometry &
Bonding Theories
Overview


Molecular Shapes
VSEPR Model
 Predicting Shapes
 Effect of Nonbonding Electrons


Polarity of Molecules
Covalent Bonding

Hybrid Orbitals
 sp, sp2, sp3
 hybrids containing d orbitals

Multiple Bonds
 sigma (s) & pi (p)
 localized & delocalized

Molecular Orbitals
 electron configurations & bond order
 diamagnetism & paramagnetism
Molecular Shapes & VSEPR

Shapes defined by bond angles
 linear, 180° angles
 trigonal planar, 120° angles
 tetrahedral, 109.5° angles

VSEPR
 Valence Shell Electron Pair Repulsion theory
 electron pairs are arranged symmetrically with
maximum separation

Two electron pairs
 180° apart
 linear geometry
180°
••
••

Three electron pairs
 120° apart
 trigonal planar geometry
••
••
••
120°

Four electron pairs
 109.5° apart
 tetrahedral geometry
••
109.5°
••
••
••
Five electron pairs
 angles of 90° and 120°
 trigonal bipyramidal (TBP) geometry
••
120°
••

90°
Six electron pairs
 angles of 90°
 octahedral geometry
••
••
••

90°
Geometries

Electron pair geometry
 arrangement of electron pairs around a central
atom

Molecular Geometry
 arrangement of atoms around a central atom

When all electron pairs are bonding pairs
 electron pair geometry = molecular geometry

When there are unshared electron pairs
 electron pair geometry  molecular geometry

To determine electron pair geometry
 draw Lewis dot structure
 count shared & unshared electron pairs around
central atom
• a multiple bond is counted as only one bonding pair when
predicting geometry
 determine electron pair geometry based on the
number of electron pairs
•
•
•
•
•
2 pair =
3 pair =
4 pair =
5 pair =
6 pair =
linear
trigonal planar
tetrahedral
trigonal bipyramidal
octahedral
Molecular Geometries with One or More
Unshared Pairs

Two Pairs
 electron pair geometry
linear
 bonding pairs
2
 non-bonding pairs
0
 molecular geometry
linear

Two electron pairs
molecular geometry
electron pair geometry
180°
••
••

Three Pairs
 electron pair geometry
trigonal planar
 bonding pairs
3
2
 non-bonding pairs
0
1
 molecular geometry
trig. pl.
bent

Three electron pairs
molecular geometry
electron pair geometry
trigonal planar
trigonal planar
••
••
••
120°
bent
••

Four Pairs
 electron pair geometry
tetrahedral
 bonding pairs
4
3
2
 non-bonding pairs
0
1
2
 molecular geometry
tet.
trig. pyr. bent

Four electron pairs
tetrahedral
molecular
geometry
electron pair
geometry
••
trigonal pyramid
••
109.5°
••
bent
••
••
••
••

Five Pairs
 electron pair geometry
trigonal bipyramid
 bonding pairs
5
4
3
2
 non-bonding pairs
0
1
2
3
 molecular geometry
tbp
seesaw T-shp. Lin.
molecular geometry
TBP
electron pair
geometry
seesaw
120°
90°
••
Five electron pair
••
••
••
••
••
T-shaped
••
linear
••


Six Pairs
 electron pair geometry
octahedral
 bonding pairs
6
5
4
2
 non-bonding pairs
0
1
2
4
 molecular geometry
oct
sq.pyr. sq. pl. lin.
Six electron pairs
molecular geometry
••
••
••
octahedral
••
••
square
planar
••
••
square
pyramid
••
••
••

electron pair geometry
linear
90°
Molecular Polarity

Molecules are always non-polar if all covalent
bonds are non-polar
 N2, P4, Cl2

Molecules with polar bonds can be polar or nonpolar
 H - Cl polar bond, polar molecular
 O=C=O two polar bonds but total molecule is
non-polar
d+
H
Cl
d-
equal but opposite
forces cancel out 
non-polar molecule
d+
d-
O
C
d-
O
are these dipole moments equal & opposite?
d-
O
d+
H
H d+
is this molecule polar?
yes
no
are these bond dipole moments equal & opposite?
d-
Cl
Cl
d-
yes
C
d+
d-
Cl
Cl d-
is this molecule polar?
no
are these bond dipole moments equal & opposite?
d+
H
H
d+
no
C
d-
Cl
Cl d-
is this molecule polar?
yes
Single and Multiple Bonds

s (sigma) bonds
 always the first bond between two atoms
 single bonds are localized between two atoms
• orbitals from two atoms overlap, allowing electrons to be shared
• electron density is on the internuclear axis
••
C
localized electrons
C

p (pi) bonds
 the second & third bonds between two atoms
 p bond electrons can be delocalized over
several atoms to form resonance structures
• electron density is above & below the internuclear axis
electron density can move
or delocalize
••
••
electron
density above
& below--p
bond
C C C C
C C C C
internuclear axis
Hybridization


allows for greater number of bonds
types of hybridization
 sp mixing of one s orbital & one p orbital
• 
2s
2p
sp
 sp2 mixing of one s orbital &
• 
2s
2p
sp2
 sp3 mixing of one s orbital &
• 
2s
2p
p
two p orbitals
p
three p orbitals
sp3

in sp hybridization
 the two sp hybrid orbitals form two s bonds
with linear geometry
 remaining two p orbitals form p bonds

in sp2 hybridization
 the three hybrid orbitals form three s bonds
with trigonal planar geometry
 the remaining one p orbital forms a p bond

in sp3 hybridization
 the four hybrid orbitals form four s bonds
with tetrahedral geometry
 sp3 hybrid atoms can form no p bonds as they
have no unhybridized p orbitals
Molecular Orbitals

mathematical combinations of atomic orbitals
 delocalized over whole molecule
 n atomic orbitals produce n molecular orbitals
• ½ are bonding orbitals and ½ are antibonding orbitals

bond order
 # bonding electrons - # antibonding electron
2

electron configuration of diatomic, homonuclear
molecules
s*
p*
p
s
s*
s
MO’s from p orbital
combination
MO’s from s orbital
combination

electron configuration of diatomic, homonuclear
molecules with interaction of the 2s and 2p
orbitals
s*
p*
s
p
s*
s
relative positions
switched
B.O. = 1
s*
s
B.O. = 3
s*
p*
s
p
s*
s
H2
2 electrons
p*
p
s*
s
N2
10 electrons
B.O. = 0
s*
s
B.O. = 2
s*
p*
p
s
s*
s
He2
4 electrons
p*
p
s*
s
O2
12 electrons