Atomic Structure 1.2

Molecular Orbital Theory (MOT)

In 1932, Mulliken developed the Molecular Orbital (MO) theory to explain:

• Formation of chemical bonds
• Relative bond strengths
• Paramagnetic or diamagnetic nature of molecules

Each molecular orbital arises from interactions between orbitals of atomic centres in the molecule, and such interactions are:

• Allowed if the symmetries of the atomic orbitals are compatible
• Efficient if the region of overlap between the two atomic orbitals is significant
• Efficient if the atomic orbitals are relatively close in energy

An important rule of MO theory is that the number of MOs formed must equal the number of atomic orbitals of the constituent atoms.

The maximum capacity of a MO is two electrons with opposite spins. Only those atomic orbitals with comparable energy and proper orientation can combine to form MOs.

The bonding MOs are represented by σ, π, etc., whereas the antibonding MOs are represented by σ^*, π^*, etc.

Inner orbitals of MOs, which do not take part in bond formation, are called non-bonding MOs.

The filling of MOs takes place according to:

• Aufbau principle — filling of MOs in order of increasing energy.
• Pauli exclusion principle — a maximum of two electrons in a MO, with opposite spins.
• Hund’s rule of maximum multiplicity — pairing of electrons in degenerate MOs does not occur until each MO has one electron.

If a molecule contains one or more unpaired electrons in its MOs, it is paramagnetic; otherwise, it is diamagnetic.