In CHEM 151, you learned about a theory of chemical bonding called valence bond theory. In this theory, chemical bonds (which contained valence electrons being shared between two bonding atoms) are formed when orbitals from each of the bonding atoms containing each atom’s valence electrons overlap.
Molecular Orbital Theory uses a different philosophy. The orbitals you learned about in CHEM 151 are called atomic orbitals – as each orbital belongs to the atom. In molecular orbital theory, once atoms are bonded into molecules the orbitals from each atom are no longer valid — instead the molecule has orbitals belonging to the entire molecule.
Neither theory is really incorrect — they are just two different ways of explaining the location of the electrons in a molecule (kind of like saying “is a zebra black with white stripes or white with black stripes”). Molecular orbital theory does handle some aspects well that valence bond theory does not. If explains why some molecules, such as O2, are paramagnetic (review the Bruce’s notes on Transition Metals and Coordination Chemistry if you don’t remember what paramagetic means). It also explains odd electron molecules such as NO better than valence bond theory.
We will begin our brief look at molecular orbital theory with a general overview before looking at examples in simple molecules including diatomic molecules involving elements from the 1st and second rows of the periodic table.
