Molecular Orbital Theory II

A molecular orbital diagram represents the relative energies of the constituent atomic orbitals, shown on the left and right, and the resultant molecular orbitals, shown in the middle.

As with atomic orbitals, each molecular orbital can accommodate a maximum of two electrons with opposite spins.The electrons from the overlapping atomic orbitals are placed in the molecular orbitals from lowest to highest energy, following Hund’s rule as needed.

For example, the two electrons of molecular hydrogen are both in the lower-energy σ_1s bonding molecular orbital.

Molecular orbital theory predicts the stability of covalent bonds from the bond order of the molecule, which is the number of electrons in bonding orbitals minus the number of electrons in antibonding orbitals divided by two.

A bond order of greater than zero indicates that one or more covalent bonds can exist, whereas a bond order of zero means that bonds should not exist. Molecular hydrogen has a bond order of one and a single bond between the hydrogens. However, covalently bound dihelium does not exist as it would have a bond order of zero.

Molecular orbital diagrams typically only include valence orbitals because the contributions of valence electrons to chemical bonding are more significant than that of core electrons.

For example, the 1s–1s overlap in dilithium has no net contribution to the bond order of the molecule. The valence electrons in the σ_2s bonding molecular orbital are responsible for its bond order of 1.

Diatomic molecules of period 2 must consider both 2s and 2p orbitals. The 2p atomic orbitals are higher in energy than the 2s orbitals, so the corresponding molecular orbitals follow the same trend.

The amount of overlap between the atomic orbitals determines how much more or less stable the molecular orbitals are. Thus, the π orbitals formed from 2p–2p sideways overlap typically fall between the σ orbitals formed from 2p–2p end-to-end overlap along the internuclear axis.

However, effects like orbital mixing between 2s and 2p orbitals, which occurs in diatomic boron, carbon, and nitrogen, leads to a different order, with the σ_2p orbital higher in energy than the π_2p.

In heteronuclear diatomic molecules, such as lithium hydride, the more electronegative atom will have lower-energy atomic orbitals. The resulting bonding molecular orbital energies will be closer to the more electronegative atomic orbital whereas the antibonding molecular orbital energies will be closer to the less electronegative atomic orbital.

Molecular orbital theory is also useful for polyatomic molecules like benzene. The Lewis model of benzene cannot accurately represent its delocalized electrons, whereas molecular orbital theory assigns those electrons to three π bonding molecular orbitals covering the entire carbon ring.