Orbital Hybridization and Molecular
Orbitals
Chapter 9
Historical Models
G.N.Lewis and I. Langmuir
- laid out foundations for molecular structures
- Ionic species were formed by electron transfer
-Covalent molecules arise from electron sharing
 cannot predict molecular geometries?
Valence shell electron pair repulsion theory (VSEPR)
- predicts molecular shapes based on valence electrons
- uses the Lewis dot structures and electron repulsions to construct models
 how do orbitals allow these geometries?
Valence bond theory (VB)
- a molecule arises from interaction (overlap) of complete atoms
- binding through localized overlap of half filled valence-shell atomic orbitals retaining
their original character (s, p, d, f)
 could not explain experimentally determined molecular properties?
Molecular orbital theory (MO)- a more complete theory
- a molecule is formed by the overlap of atomic orbitals to form molecular orbitals
- electrons are then distributed into MOs
-A molecule is a collection of nuclei with the orbitals delocalized (spread out) over the
entire molecule
Valence Bond Theory
Sigma Bond Formation
Two s
orbitals
overlap
Two p
orbitals
overlap
Valence Bond (VB) Theory
Covalent bonds are formed by the overlap of
atomic orbitals
-Atomic orbitals on the central atom can mix and
exchange their character with other atomic
orbitals in a molecule
-This process is called hybridization
-Hybrid Orbitals: create geometries with the
same shapes as seen in VSEPR theory
Valence Bond Theory
Linus Pauling
-valence electrons are
localized between atoms (or
as lone pairs)
-half-filled atomic orbitals
overlap to form bonds
Valence Bond (VB) Theory
Regions of High
Electron Density
(“things” attached)
2
3
Electronic
Geometry
Hybridization
Linear
Trigonal planar
sp
sp2
4
5
Tetrahedral
Trigonal
bipyramidal
sp3
sp3d
6
Octahedral
sp3d2
Molecular Shapes and Bonding
In the next sections we will use the following
terminology:
A = central atom
B = bonding pairs around central atom
U = lone pairs around central atom
Linear Electronic Geometry: AB2
Some examples of molecules with this
geometry are:
BeCl2, BeBr2, BeI2, HgCl2, CdCl2
-All of these examples are linear,
nonpolar molecules
-Important exceptions occur when the
two substituents are not the same
BeClBr or BeIBr will be linear and polar
Linear Electronic Geometry: AB2
1-p
1-p
Trigonal Planar Electronic Geometry:
AB3
Some examples of molecules with this geometry
are:
BF3, BCl3
-All of these examples are trigonal planar,
nonpolar molecules
-Important exceptions occur when the three
substituents are not the same
BF2Cl or BCI2Br will be trigonal planar and polar!
Trigonal Planar Electronic Geometry:
AB3
Trigonal Planar
How to account for 3 bonds 120o apart using a spherical s orbital
and p orbitals that are 90o apart?
-Pauling said to modify VB approach with ORBITAL
HYBRIDIZATION
-mix bonding orbitals to form a new set of orbitals
-HYBRID ORBITALS : give the maximum overlap in the
correct geometry
Valance Bond Theory
2p
2s
hydridize orbs.
2
rearrange electrons
three sp
hybrid orbitals
unused p
orbital
Valence Bond Theory
The three hybrid orbitals are made from 1-s
orbital and 2-p orbitals  3-sp2 hybrids
Now we have 3, half-filled HYBRID orbitals that can
be used to form B-F sigma bonds
Valance Bond Theory
Bonding in BF3
•• ••
F ••
••••
F
••
B
Boron configuration

•••
F• 1s
••

2s

2p
planar triangle
angle = 120o
Tetrahedral Electronic Geometry: AB4
Some examples of molecules with this geometry are:
CH4, CF4, CCl4, SiH4, SiF4
-All of these examples are tetrahedral, nonpolar
molecules
-Important exceptions occur when the four
substituents are not the same
CF3Cl or CH2CI2 will be tetrahedral and polar
Tetrahedral Electronic Geometry: AB4
What’s the electron configuration of carbon?
How many free electrons does it have?
How many bonds can it form?
Tetrahedral Bonding
How do we account for 4 C—H
sigma bonds 109o apart?
-Need to use 4 atomic
orbitals
-s, px, py, and pz to form
4 new hybrid orbitals
-hybrid orbitals point in the
correct orientation
109o
Tetrahedral Electronic Geometry: AB4
Tetrahedral Electronic Geometry: AB4
sp3
Tetrahedral Electronic Geometry: AB3U
Some examples of molecules with this geometry
are:
NH3, NF3, PH3, PCl3, AsH3
-These molecules are examples of central
atoms with lone pairs of electrons
-Thus, the electronic and molecular geometries are
different
-All three substituents are the same but molecule
is polar
-NH3 and NF3 are trigonal pyramidal, polar
molecules
Steps in predicting the hybrid orbitals used by an atom in bonding:
1. Draw the Lewis structure
2. Determine the electron pair geometry using the VSEPR model
3. Specify the hybrid orbitals needed to accommodate the electron pairs in the
geometric arrangement
NH3
1. Lewis structure
2. VSEPR indicates tetrahedral geometry
with one non-bonding pair of electrons
(structure itself will be trigonal pyramidal)
3. Tetrahedral arrangement indicates four
equivalent electron orbitals
Tetrahedral Electronic Geometry:
AB2U2
Some examples of molecules with this geometry
are:
H2O, OF2, H2S
-These molecules are examples of central atoms
with two lone pairs of electrons
-Thus, the electronic and molecular geometries are
different
-Both substituents are the same but molecule is
polar.
-Molecules are angular, bent, or V-shaped and
polar
Orbital Hybridization
Bonds
Shape
(“Things”)
Hybrid
orbital
Remaining
orbitals
2
linear
sp
2 p’s
3
trigonal
planar
sp2
1p
4
tetrahedral sp3
none
Trigonal Bipyramidal Electronic
Geometry: AB5, AB4U, AB3U2, and
AB2U3
Some examples of molecules with this geometry
are:
PF5, AsF5, PCl5, etc.
-These molecules are examples of central atoms with
five bonding pairs of electrons
-The electronic and molecular geometries are the same
-Molecules are trigonal bipyramidal and nonpolar
when all five substituents are the same
-If the five substituents are not the same polar molecules
can result, AsF4Cl is an example.
Trigonal Bipyramidal Electronic
Geometry: AB5, AB4U, AB3U2, and
AB2U3
Octahedral Electronic Geometry: AB6,
AB5U, and AB4U2
Some examples of molecules with this geometry are:
SF6, SeF6, SCl6, etc.
-These molecules are examples of central atoms with
six bonding pairs of electrons
-Molecules are octahedral and nonpolar when all six
substituents are the same
-If the six substituents are not the same polar molecules
can result, SF5Cl is an example.
Octahedral Electronic Geometry: AB6,
AB5U, and AB4U2
Molecular Geometry
H
H C
H
H
Octahedral Electronic Geometry: AB6,
AB5U, and AB4U2
Polarity
Molecular Geometry
H
H C
H
H
Carbon-Carbon Double Bonds
C atom has four valence electrons
-Three electrons from each C atom are in sp2 hybrids
-One electron in each C atom remains in an unhybridized p
orbital
2s 2p
three sp2 hybrids 2p
C    


Double Bonds
The single 2p orbital is perpendicular to the
trigonal planar sp2 lobes (containing 3 e-)
-The fourth electron is in the p orbital
Side view of sp2 hybrid with p
orbital included
Double Bonds
An sp2 hybridized C atom has this shape
-Remember there will be one electron in
each of the three lobes
Top view of an sp2 hybrid
Double Bonds
The portion of the double bond formed
from the head-on overlap of the sp2
hybrids is designated as a s bond
p
p
Triple Bonds
Ethyne or acetylene, C2H2, is the simplest triple
bond containing organic compound
-Compound must have a triple bond to obey
octet rule
Triple Bonds
Lewis Dot Formula
H ·· C ·· ·· ·· C ·· H
or
H C C H
VSEPR suggests regions of high electron
density are 180o apart
Carbon-Carbon Triple Bonds
Carbon has 4 electrons
-Two of the electrons are in sp hybrids
-Two electrons remain in unhybridized p
orbitals
C [He]
2s

2p
2-sp hybrids
 

H-C=C-H
2p

Triple Bonds
A s bond results from the head-on overlap
of two sp hybrid orbitals
py
pz
Triple Bonds
The unhybridized p orbitals form two p bonds
-Note that a triple bond consists of one s and two
p bonds
- py
- pz
Summary of Geometries
Molecular Orbital (MO) Theory
VSEPR and VB theory are good to explain the molecular
shape
-BUT they did not explain the magnetic or spectral properties
of molecules
-Molecular orbital theory does!
MOs are derived from the addition and subtraction of
atomic orbitals represented as wavefunctions to form
molecular orbitals
There are two possible combinations:
-Adding two atomic orbitals forms a bonding MO
-Subtracting two atomic orbitals forms an antibonding MO
Note: The number of atomic orbitals contributed equals
the number of molecular orbitals generated
Molecular Orbitals
Consider H2
Molecular Orbital
-Bonding MO is lower in energy than the unbound
atomic orbitals
-Antibonding MO is higher in energy
Molecular Orbital Diagram
Electrons fill according to Hund’s rule
Molecular Orbital Theory
Bond Order (B.O.)
B.O. = 1/2 (B - AB)
B = bonding electrons AB = antibonding electrons
Molecular Orbital Theory
Example: He2
MO Diagram for He2+ and H2-
AO of
He
s*1s
AO of
He+
s1s
MO of
He2+
He2+ bond order = ??
Energy
s*1s
AO of
AO of
H
Hs1s
MO of
H2 -
H2- bond order = ??
MO Diagram for He2+ and H2-
s*1s
1s
1s
MO of
He2+
1s
1s
s1s
s1s
AO of
He
Energy
s*1s
AO of
He+
He2+ bond order = 1/2
AO of
H-
MO of
H2 -
AO of
H
H2- bond order = 1/2
Orbital Interaction for Li2 Molecule
Li atom - 1s22s1
s*2s
Li 2s
Li 2s
s2s
Bond order for Li2?
s*1s
Li 1s
Li 1s
s1s
Be2?
Li2+?
p Orbital Bonding
p Orbital Bonding
O2 molecule is an example with sigma and pi bonds forming
between atoms
-MO theory predicts that oxygen will be paramagnetic
½ ( 8 - 4) = 2