Chapter 10: Covalent Bond Theories
Valence Bond Theory
Molecular Orbital Theory
First a couple of jokes… (You’ll need a laugh this chapter!)
If a bear from Yosemite and a bear from
Alaska fall into the water, which one dissolves
faster?
Which is lower energy, hamburger or steak?
My name is Bond - Ionic Bond; taken not
Valence Bond Theory –
Covalent Bonds form from the __________ of
atomic orbitals.
• Consider a BeF2 molecule:
F
Be
F
Valence Bond Theory
F
F
1s22s22p5
One unpaired eBe
1s22s2
No unpaired electrons
1s22s22p5
One unpaired eHow can Be
form covalent
bonds with F if
it doesn’t have
unpaired
electrons???
Valence Bond Theory
Promote one of the 2s e- to a 2p orbital
Be
no unpaired electrons
1s22s2
Be
2 unpaired electrons
1s22s12p1
Valence Bond Theory
F
1s22s22p5
One unpaired eBe
1s22s12p1
2 unpaired electrons
F
1s22s22p5
One unpaired ePredicted overlaps:
2s (Be) – 2p (F)
2p (Be) – 2 p (F)
Implies two different
kinds of Be – F bonds!
Valence Bond Theory
• Both of the Be – F bonds in BeF2 are
identical!
• The solution:
– Mix the Be 2s orbital with one of the Be 2p
orbitals to form two hybrid orbitals
• atomic orbitals formed by mixing 2 or more atomic
orbitals on an atom
Valence Bond Theory
F
F
1s22s22p5
1s22s22p5
Be
1s
2sp
2p
• When Be forms covalent bonds with two F,
each sp hybrid orbital on the Be atom
overlaps with a p orbital located on a F
atom.
Valence Bond Theory
• The BeF2 molecule is linear.
• sp hybridization implies that the electron
geometry around the central atom is linear.
The sp2 hybrid orbitals in BF3.
The sp3 hybrid orbitals in CH4.
The sp3 hybrid orbitals in NH3.
The sp3 hybrid orbitals in H2O.
The sp3d hybrid orbitals in PCl5.
The sp3d2 hybrid orbitals in SF6.
Show the orbital hybridization of SF4
Multiple Bonds
All of the bonds shown thus far have been single bonds which result in sigma bonds
Sigma bond (s)  A bond where the line of electron density
is concentrated symmetrically along the line connecting the
two atoms. Results from head to head overlap of orbitals
Pi bond (p)  A bond where the overlapping regions exist
above and below the internuclear axis.
Pi bonds result from
side to side overlap of
atomic orbitals. Pi
bonds cannot form from
hybrid orbitals
A double bond consists of one
sigma (hybrid) and one pure pi
bond.
A triple bond consists of one
hybrid sigma and 2 pure pi
bonds
Example: H2C=CH2
Example: H2C=CH2
Only sigma bond pairs influence molecular
geometry. What is the geometry around each
central atom?
Multiple Bonding:
Identify sigma and pi bonds in the following molecules and
name the orbital hybridization involved.
• Formaldehyde: H2CO
• Acetylene: C2H2
• Nitrogen gas
• Acetonitrile, CH3CN
The paramagnetic
properties of O2
Using the valence
bond theory,
explain the
paramagnetism of
oxygen gas.
Molecular Orbital Theory –
Another theory for the formation of covalent bonds which
explains phenomena for which the Valence Bond Theory
can’t account .
• When 2 atomic orbitals overlap, two new
molecular orbitals are formed:
– Bonding molecular orbital
– Antibonding molecular orbital
Electrons are spread or delocalized over the
whole molecule instead of remaining in
atomic orbitals
Molecular Orbital Theory
• Bonding molecular orbital
– Addition of overlapping atomic orbitals results
in an area of increased e- density; increased
probability e- will be found here.
– lower energy than atomic orbital
Bonding MO
Molecular Orbital Theory
• Antibonding molecular orbital
– An area where the probability of finding an
electron is reduced
• Results from subtracting one atomic orbital from
the other
– higher energy than atomic orbital
Antibonding
MO
Bonding
MO
Molecular Orbital (MO) Theory
ANTBONDING
These two new orbitals have
different energies.
BONDING
The one that is lower in energy is called the bonding orbital,
The one higher in energy is called an antibonding orbital.
Molecular Orbital Theory
• The MO diagram for H2 molecule:
s*1s
1s
H atom
Placing e-s in MO’s follows
the Aufbau principle, Pauli’s
and Hund’s Rule.
1s
s1s
H atom
bonding electrons
H2 molecule
Molecular Orbital Theory
• The MO diagram for He2 molecule:
s*1s
antibonding electrons
1s
He atom
1s
s1s
He atom
bonding electrons
He2 molecule
Adding electrons to antibonding molecular orbitals destabilizes the
molecule.
Molecular Orbital Theory
• In MO Theory, the stability of a covalent bond
can be related to its bond order:
• Bond order = (# bonding e- - # antibonding e-)
2
– Old def.: # e- pairs / # links formed
– Single bond: bond order = 1
– Double bond: bond order = 2
– Triple bond: bond order = 3
– Fractional bond orders also exist! (e.g. He2+1)
Energy level diagrams / molecular
orbital diagrams
Molecular Orbital Theory
• The bond order for H2 molecule:
Bond Order = (2 - 0) = 1
2
• The hydrogen atoms in an H2 molecule
are held together by a single bond.
Molecular Orbital Theory
• The bond order for an He2 molecule:
Bond order = (2 - 2) = 0
2
• No bond exists between two He atoms.
Example: Li2
Atomic orbitals combine to form molecular orbitals best when the atomic orbitals
are of equal energy.
Second-Row Diatomic Molecules
Electron Configurations for B2 through Ne2
Relative MO energy levels for Period 2 homonuclear diatomic molecules.
without 2s-2p
mixing
with 2s-2p
mixing
MO energy levels
for O2, F2, and Ne2
MO energy levels
for B2, C2, and N2
For B2, C2, and N2 the interaction is so strong that the s2p is
pushed higher in energy than p2p orbitals
Molecular Orbital Theory
Example: Complete the MO diagram for N2.
Calculate the bond order.
# valence electrons for each N atom =
Total # of valence electrons =
Molecular Orbital Theory
s*2p
p*2p
# bonding e- =
s2p
# antibonding e- =
p2p
Bond order =
s*2s
s2s
Molecular Orbital Theory
Example: Complete the MO diagram for the
O2 molecule. Determine the bond order.
# valence electrons for each O atom =
Total # of valence electrons =
Molecular Orbital Theory
s*2p
p*2p
# bonding e- =
p2p
# antibonding e- =
s2p
Bond order =
s*2s
s2s
Isomers
1,1-dichloroethylene
1,2-dichloroethylene
cis-1,2-dichloroethylene
trans-1,2-dichloroethylene