MOLECULAR ORBITALS

When atoms combine into a molecule, the electrons contributed to the chemical bonds by each atom are no longer localized on individual atoms, but “belong” to the entire molecule. Consequently, atomic orbitals are no longer appropriate descriptions for the state of electrons in molecules. Instead, molecular orbitals (MO), which are orbitals for the entire molecule, are used. In Molecular Orbital Theory we describe electrons in overlapping orbitals of different atoms as being localized in the bonds between the two atoms involved. We then use hybridization to help account for the geometry of a molecule.

Quantum theory gives us a few simple rules that allow to derive these molecular orbitals without any calculations: 1. The combination of 2 atomic orbitals (AOs) gives 2 molecular orbitals (MOs). 2. One is derived by addition of the two atomic orbitals (in phase) in the region of overlap (bonding MO) 3. The other molecular orbital is derived by subtraction of the two atomic orbitals (out of phase) in the region of overlap (antibonding molecular orbital, or antibonding MO). Subtraction of two atomic orbitals is the same as reversing the sign of either one and adding. 4. The two MOs have different energies. One has a lower energy than the starting atomic orbitals (bonding MO), and the other has a higher energy (antibonding MO). 5. The valence electrons from the starting atoms are distributed in the new molecular orbitals in accord with the Pauli principle and Hund’s rules.

MOLECULAR ORBITALS DIAGRAMS: example H2 MOLECULE

------MO DIAGRAMS FOR DIATOMIC OMONUCLEAR

There are two possibilities: *diagram 1: No mixing between 2s and 2p orbitals: occurs when s and p atomic orbitals are far in energy (≥1/2 filled 2p orbitals)

** diagram 2 Mixing between s and p orbitals (sp mixing): occurs when the s and p atomic orbitals are close in energy ( ≤ 1/2 filled 2p orbitals)

**diagram 2