Problems With Valence Bond Theory
VSEPR Theory - electron pair
repulsions influence molecular shape
Valence Bond Theory - atoms form
bonds by overlapping hybrid orbitals
Applied to O2
O=O
There’s a problem with this Lewis structure for oxygen,
however, and it has to do with magnetism in atoms -
Cause of Magnetism = ELECTRON SPIN
Diamagnetic materials = all electron
spins are “paired”, so the magnetic fields
cancel out
N
S
S
N
=
Zn = [Ar]3d104s2
Paramagnetic materials = unpaired
electrons so there is a net magnetism
Mn = [Ar]3d54s2
N
=
S
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In the Lewis structure for
oxygen, all of the
electrons are paired up, so
the molecule should be
diamagnetic, but
experiments prove that it
is PARAMAGNETIC.
An additional refinement
in bonding theory is
necessary =
Molecular Orbitals - Preliminary Ideas
Don’t forget that electrons behave like WAVES, and
there are WAVE FUNCTIONS () that describe the
electron position in space = ATOMIC ORBITALS (2)
e'
2
Waves (electrons) can interfere with each other,
either CONSTRUCTIVELY or DESTRUCTIVELY
Principles of Molecular Orbital Theory
1. The total number of molecular orbitals = total
number of atomic orbitals contributed by the
bonding atoms
2. Bonding MO’s are lower in energy (more stable)
than antibonding MO’s
3. Electrons occupy molecular orbitals following the
Pauli Exclusion Principle (spins pair up) and
Hund’s Rule (remain unpaired as long as an
empty orbital is available of the same energy)
3
Types of Molecular Orbitals
Bonding and Antibonding
-bond formation involving p-orbitals
*2p
2p
4
-bond formation involving p-orbitals
*2p
2p
*2p
2p
Molecular Orbital Energy Levels
Atomic
orbitals
*2p
*2p
2p
2p
2p
2s
2p
Atomic
orbitals
2s
*2s
1s
2s
1s
*1s
1s
5
Molecular Orbital Diagram: H2
The two 1s orbitals may be added or subtracted
to yield two sigma MOs (1 bonding/1 antibonding).
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Bond Order and Stability
Bond Order = 1/2 (number or bonding e−
− number or antibonding e−).
Bond order in H2 = 1
(Stable).
Bond order in He2 = 0
(Not Stable).
MO Diagrams for N2 and O2
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MO Scheme for N2
 Electron configuration
for N2 =
 Bond order =
• N2 has three bonds.
• N2 has no unpaired
electrons.
MO Scheme for O2
 Electron configuration for
O2 =
 Bond order=
• O2 has two bonds
• O2 has two unpaired
electrons in π2p*
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Other Diatomic Molecules
MO Theory correctly predicts both the magnetic properties
and bond orders of diatomic molecules.
MO Theory: Summary
 Advantages:
• Provides the most complete picture of
covalent bonding, including bond types
and bond orders.
• Accounts for magnetic properties.
 Disadvantage:
• The most difficult to apply to large
molecules; does not account for molecular
shape.
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