Molecular orbital theory

Atomic orbitals of the bonded atoms combine into molecular orbitals that belong to the whole molecule; the number of MOs formed equals the number of atomic orbitals combined.

MOs form by linear combination of atomic orbitals (LCAO): addition gives a bonding MO of lower energy, subtraction gives an antibonding MO (marked *) of higher energy.

Combining orbitals must have similar energy, the same symmetry about the molecular axis, and must overlap well.

MOs are filled following the aufbau principle, Pauli's exclusion principle and Hund's rule.

Energy order up to N₂: σ1s < σ*1s < σ2s < σ*2s < (π2px = π2py) < σ2pz < (π*2px = π*2py) < σ*2pz.

For O₂ and F₂ the σ2pz orbital lies below the two π2p orbitals.

Bond order = ½(Nb − Na); a positive value means a stable molecule, zero means the molecule does not exist, and higher bond order means shorter and stronger bonds.

Unpaired electrons in MOs make a species paramagnetic; if all electrons are paired it is diamagnetic.