As described in Part 2, the filling up of molecular orbitals follow the same principles and rules as filling up atomic orbitals. Now, we can use energy level diagrams to rationalise differences in the stability of molecules. Let us look at the molecules H­₂⁺, H₂, He₂⁺ and He₂

The molecular orbitals would look like:

The table below shows the electronic configurations, dissociation energies and bond lengths of the respective molecules.

MO theory helps us explain the difference in the dissociation energies of the molecules. H­₂⁺, H₂ and He₂⁺ are stable with respect to their dissociation into their constituents because they have more electrons in the bonding MO than the anti-bonding MO. He₂ however has two electrons in the bonding MO and two in the anti-bonding MO. The anti-bonding MO is raised more in energy than the bonding MO is lowered, so we can expect the two anti-bonding electrons to outweigh the two bonding ones. As a result, He₂ is unstable and will dissociate into He atoms.

We have discussed the concept of bond order in the previous post. The bond order can also be considered when determining whether a bond forms (i.e. whether the molecule is stable). For example,
H­₂⁺ has 1e in bonding, 0 in anti-bonding –> bond order = 1/2
H₂ has 2e in bonding, 0 in anti-bonding –> bond order = 1
He₂⁺ has 2e in bonding, 1 in anti-bonding –> bond order = 1/2
However He₂ has 2e in bonding, 2 in anti-bonding –> bond order = 0. Therefore it is unstable