The Diatomic Molecules of the Second Period

Eight possible homonuclear diatomic molecules might be formed by the atoms of the second period of the periodic table: Li2, Be2, B2, C2, N2, O2, F2, and Ne2. However, we can predict that the Be2 molecule and the Ne2 molecule would not be stable. We can see this by a consideration of the molecular electron configurations.

We predict valence molecular orbital electron configurations just as we predict electron configurations of atoms. Valence electrons are assigned to valence molecular orbitals with the lowest possible energies. Consistent with Hund’s rule, whenever there are two or more degenerate molecular orbitals, electrons fill each orbital of that type singly before any pairing of electrons takes place.

As we saw in valence bond theory, σ bonds are generally more stable than π bonds formed from degenerate atomic orbitals. Similarly, in molecular orbital theory, σ orbitals are usually more stable than π orbitals. However, this is not always the case. The MOs for the valence orbitals of the second period are shown in the figure below. Looking at Ne2 molecular orbitals, we see that the order is consistent with the generic diagram shown in the previous section. However, for atoms with three or fewer electrons in the p orbitals (Li through N) we observe a different pattern, in which the σp orbital is higher in energy than the πp set. Obtain the molecular orbital diagram for a homonuclear diatomic ion by adding or subtracting electrons from the diagram for the neutral molecule.

This shows the MO diagrams for each homonuclear diatomic molecule in the second period. The orbital energies decrease across the period as the effective nuclear charge increases and atomic radius decreases. Between N2 and O2, the order of the orbitals changes.

You can practice labeling and filling molecular orbitals with interactive tutorials from the University of Sydney.

This switch in orbital ordering occurs because of a phenomenon called s-p mixing. s-p mixing does not create new orbitals; it merely influences the energies of the existing molecular orbitals. The σs wavefunction mathematically combines with the σp wavefunction, with the result that the σs orbital becomes more stable, and the σp orbital becomes less stable. Similarly, the antibonding orbitals also undergo s-p mixing, with the $σ_s^*$ becoming more stable and the $σ_p^*$ becoming less stable.

Without mixing, the MO pattern occurs as expected, with the σp orbital lower in energy than the πp orbitals. When s-p mixing occurs, the orbitals shift as shown, with the σp orbital higher in energy than the πp orbitals.

s-p mixing occurs when the s and p orbitals have similar energies. The energy difference between 2s and 2p orbitals in O, F, and Neis greater than that in Li, Be, B, C, and N. Because of this, O2, F2, and Ne2 exhibit negligible s-p mixing (not sufficient to change the energy ordering), and their MO diagrams follow the normal pattern, as shown in the figure above. All of the other period 2 diatomic molecules do have s-p mixing, which leads to the pattern where the σp orbital is raised above the πp set.

Using the standard MO diagrams, we can add in the electrons and determine the molecular electron configuration and bond order for each of the diatomic molecules. By filling these in, you can determine that Be2 and Ne2 molecules would have a bond order of 0, and in fact these molecules do not exist.

Electron Configuration and Bond Order for Molecular Orbitals in Homonuclear Diatomic Molecules of Period Two Elements

MoleculeElectron ConfigurationBond Order
Be2 (unstable)$(σ_{2s})^2(σ_{2s}^*)^2$0
C2 $(σ_{2s})^2(σ_{2s}^*)^2(π_{2p_y},π_{2p_z})^4$ 2
N2 $(σ_{2s})^2(σ_{2s}^*)^2(π_{2p_y},π_{2p_z})^4(σ_{2p_x})^2$3
F2$(σ_{2s})^2(σ_{2s}^*)^2(σ_{2p_x})^2(π_{2p_y},π_{2p_z})^4(π_{2p_y}^*,π_{2p_z}^*)^4$ 1
Ne2 (unstable) $(σ_{2s})^2(σ_{2s}^*)^2(σ_{2p_x})^2(π_{2p_y},π_{2p_z})^4(π_{2p_y}^*,π_{2p_z}^*)^4(σ_{2p_x}^*)^2$0

The combination of two lithium atoms to form a lithium molecule, Li2, is analogous to the formation of H2, but the atomic orbitals involved are the valence 2s orbitals. Each of the two lithium atoms has one valence electron. Hence, we have two valence electrons available for the σ2s bonding molecular orbital. Because both valence electrons would be in a bonding orbital, we would predict the Li2 molecule to be stable. The molecule is, in fact, present in appreciable concentration in lithium vapor at temperatures near the boiling point of the element. All of the other molecules in [link] with a bond order greater than zero are also known.

The O2 molecule has enough electrons to half fill the ($π_{2p_y}^*$, $π_{2p_z}^*$) level. We expect the two electrons that occupy these two degenerate orbitals to be unpaired, and this molecular electronic configuration for O2 is in accord with the fact that the oxygen molecule has two unpaired electrons (as seen in the start of this chapter). The presence of two unpaired electrons has proved to be difficult to explain using Lewis structures, but the molecular orbital theory explains it quite well. In fact, the unpaired electrons of the oxygen molecule provide a strong piece of support for the molecular orbital theory.