VSEPR cannot explain the bonding in all compounds. Molecular orbital theory can provide an explanation for more complex molecules.

Molecular orbitals form when atomic orbitals combine. The number of molecular orbitals formed is equal to the number of atomic orbitals that combine.

→ The combination of two atomic orbitals results in the formation of a bonding molecular orbital and an antibonding orbital. Each molecular orbital can hold a maximum of two electrons.

→ The bonding molecular orbital encompasses both nuclei. The attraction of the positively charged nuclei and the negatively charged electrons in the bonding molecular orbital is the basis of bonding between atoms.

→ An antibonding orbital is designated as a $\sigma$* or $\pi$* orbital.

In a non-polar covalent bond, the bonding molecular orbital is symmetrical about the midpoint between two atoms.

Polar covalent bonds result from bonding molecular orbitals that are asymmetric about the midpoint between two atoms. The atom with the greater value for electronegativity has the greater share of the bonding electrons.

Ionic compounds are an extreme case of asymmetry, with the bonding molecular orbitals being almost entirely located around just one atom, resulting in the formation of ions.

Molecular orbitals that form by end-on overlap of atomic orbitals along the axis of the covalent bond are called sigma, $\sigma$, molecular orbitals or $\sigma$ bonds.

Molecular orbitals that form by side-on overlap of parallel atomic orbitals that lie perpendicular to the axis of the covalent bond are called pi, $\pi$, molecular orbitals or $\pi$ bonds.

The electronic configuration of an isolated carbon atom cannot explain the number of bonds formed by carbon atoms in molecules. The bonding and shape of molecules of carbon can be explained by hybridisation.