Molecular Orbital Theory
Chemistry 123
Dr. Patrick Woodward
Supplemental Lecture 2
Atomic Orbitals
Think of the electrons in atomic orbitals as standing waves. The color
coding gives the phase of the wave. When the wavefunction is colored blue
the wave is positive, when it is colored red it is negative.
See: http://winter.group.shef.ac.uk/orbitron/ for a visual depiction
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Molecular Orbital (MO) Theory
MO theory is based on the premise that electrons in molecules reside in
molecular orbitals. These molecular orbitals are formed from the overlap of
atomic orbitals. MO’s can be classified as bonding, antibonding and nonbonding.
Bonding Molecular Orbital
• Constructive Interference
• Increases electron density
between nuclei
• Forms a bond
Antibonding MO
• Destructive Interference
• Decreases electron density
between nuclei (nodal plane)
• Cancels a bond
Image taken from
http://www.sparknotes.com/chemistry/bonding/molecularorbital/section1.html
MO Diagram for H2
Antibonding MO
Bonding MO
Image taken from
http://www.sparknotes.com/chemistry/bonding/molecularorbital/section1.html
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Energy level diagrams / molecular
orbital diagrams
Single (sigma) bond
between H atoms
No bond between He
atoms
Fundamental Principles of MO Theory
Molecular Orbitals (MO’s) from Atomic Orbitals (AO’s)
1. # of Molecular Orbitals = # of Atomic Orbitals
2. The number of electrons occupying the molecular orbitals is equal to the
sum of the valence electrons on the constituent atoms.
3. When filling MO’
MO’s the Pauli Exclusion Principle Applies (2 electrons per
Molecular Orbital)
4. For degenerate MO’
MO’s, Hund's rule applies (fill each with a single
electron before pairing up 2 electrons in a single orbital).
5. AO’
AO’s of similar energy combine more readily than ones of different
energy
6. The more overlap between AOs the lower the energy of the bonding
orbital they create and the higher the energy of the antibonding orbital.
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MOs from 2p atomic orbitals
sigma
antibonding
(σ∗)
pi antibonding
(π∗)
sigma bonding
(σ)
pi bonding (π)
1 sigma interaction: Through overlap of orbitals along the internuclear axis.
2 pi interactions: Through overlap of orbitals above and below (or to the sides)
of the internuclear axis (nodal plane – zero overlap along the internuclear axis).
MO Diagram for a second row diatomic molecule (i.e. O2)
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Metal d – ligand p MO’s
M dz2 – L p σ*
M dx2-y2 – L p σ*
Ligand p
orbitals
Metal d
orbital
M dxz – L p π*
M dxy – L p π*
M dyz – L p π*
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