Molecular Orbital Theory (MOT)

CompleteFundamentals of Chemistry

MOT treats electrons as belonging to the whole molecule via molecular orbitals; it explains O₂'s paramagnetism, bond order and magnetic character.

Full notes

Definition: MOT: atomic orbitals combine to form molecular orbitals that extend over the entire molecule; electrons fill these orbitals according to Aufbau, Pauli and Hund's rules.

1. Bonding vs antibonding orbitals

Two AOs of same symmetry combine to give one bonding MO (lower energy, electron density between nuclei) and one antibonding MO σ*/π* (higher energy, node between nuclei). The number of MOs equals the number of AOs combined.

2. Energy order and filling

For second-period homonuclear diatomics (Li₂ to N₂): σ1s < σ*1s < σ2s < σ*2s < σ2pz < π2px=π2py < π*2px=π*2py < σ*2pz. From O₂ onwards, σ2pz drops below π2px/π2py because of s-p mixing effects. Electrons fill in by Hund's rule.

3. Bond order and magnetic behaviour

Bond order = (bonding − antibonding)/2. A fractional or non-integer order plus unpaired electrons explains properties.

Molecule	Bond order	Magnetic nature
B₂	1	Paramagnetic (2 unpaired)
C₂	2	Diamagnetic
N₂	3	Diamagnetic
O₂	2	Paramagnetic (2 unpaired)
F₂	1	Diamagnetic
CO	3	Diamagnetic (isoelectronic N₂)
NO	2.5	Paramagnetic

Note: Exam point: 'why is O₂ paramagnetic?' is a guaranteed question — answer with the two unpaired electrons in π* orbitals, which VBT could not explain.

Exam notes

MO = bonding (lower) + antibonding (higher). Bond order = (Nb − Na)/2. Unpaired electrons ⇒ paramagnetism. Recall O₂, N₂, NO, CO, B₂, C₂, F₂.

Important points

MOT explains paramagnetism, bond energy trends and why He₂ and Be₂ do not exist.

Common mistakes

Using the same orbital order for all second-row diatomics — O₂ onwards has different σ/π order.

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