Page 294 - Instant notes

Physical chemistry 280

Homonuclear second row diatomic molecules

Diatomic molecules composed of identical atoms N 2 and O 2, for example, are referred to
as homonuclear diatomic molecules. In the second and subsequent rows of the periodic
table, the bonding interactions involve both p and s orbital interactions. By convention,
the bonding axis is taken to be the z axis and the linear combinations of the p z orbitals
differ from those of the p x and p y orbitals.
The lobes of the atomic p z orbitals interact directly and relatively strongly along the
bonding axis to form molecular bonding and antibonding a orbital combinations. The
linear combinations of the p x and p y orbitals on the other hand give rise to two bonding
and antibonding π orbital combinations (Fig. 3). The overlap of the p x and of the p y
orbitals is smaller than that of the p z orbitals, and this is reflected in the energy of the a
bond which is lower than that of the two π bonds. The two π bonds are of equal energy
and are said to be degenerate.
Core atomic orbitals do not make a significant contribution to the bonding in a
molecule for two reasons. Firstly, there is no significant overlap of these orbitals, and the
energy of the bonding orbitals is not significantly lower than the atomic orbitals, and
secondly, each pair of bonding and antibonding is fully occupied, leaving no net bonding
contribution.
The molecular orbital diagram for a second row diatomic, O 2, is shown in Fig. 3, with
eight 2p electrons occupying the two degenerate bonding π-orbitals, the bonding σ-orbital
and two degenerate π* orbitals to give a double bond overall.
The σ bonding and antibonding molecular orbitals derived from the 1s and 2s atomic
orbitals are fully occupied and so have no net bonding effect. In oxygen, the two highest
energy electrons are placed separately into the π x* and π y* orbitals to give two unpaired
electrons, which also confer oxygen with significant paramagnetism.
As the molecular orbitals are qualitatively unchanged, the same molecular orbital
diagram may be used to describe any second row diatomic molecule or ion, by simply
entering the correct number of electrons. However, detailed analysis reveals the
importance of including all the atomic orbitals when generating the linear combination of
orbitals. The narrow energy gap between the 2s and 2p orbitals in the early part of the
second row leads to mixing of the 2s atomic orbitals with the 2p z orbitals. This raises the
energy of the 2p z generated σ orbitals, and lowers the energy of the 2s generated σ*
orbital, so that the highest occupied π and σ bonding orbitals exchange positions between
nitrogen and oxygen (Fig. 4).