The path to molecular orbital theory
The early history of the development of molecular orbital (MO) theory is reviewed. Aided by analogies to atomic spectra and based on evidence from molecular spectra in connection with the old quantum theory, a classification of electronic states of many diatomic molecules was effected early in 1926. This classification was clarified and extended with the advent of quantum mechanics, using the new concept of molecular orbitals (not so named until 1932): Hund, Mulliken. The bonding power of electrons in MOs was discussed, and with the help of the LCAO approximation, MOs were classified as bonding or antibonding. Herzberg proposed that one half the number of bonding less half the number of antibonding electrons is equal to the number of chemical bonds in not too polar diatomic molecules. As an alternative to the use of general or non-localized MOs, Hund showed the usefulness of localized and it MOs in describing the structures of single and multiple chemical bonds. The close correspondence of a pair of electrons in a localized MO to G. N. Lewis's earlier concept of an electron-pair bond is pointed out. The 'semi-united-atom' MO concept, e.g. for N2 and CO, is related to Langmuir's earlier description of the structure of these molecules. The description of the structures of polyatomic molecules using non-localized MOs, the criterion of maximum overlap in MO theory, the electronegativity scale, and the prediction of MO ionization potentials are discussed briefly. IN a discussion of 'Fifty Years of Valence Theory', I feel it is necessary to go back a few years further to the chemical theory of valence' developed by G. N. Lewis in 1916. The essential features of this theory still to a considerable extent form a useful framework for understanding the role of electrons in chemical bonding Lewis's theory spanned the previously unbridged gap between ionic valence and the kind of valence that is typical of carbon compounds. This he accomplished by using the idea of pairs of electrons shared either equally or unequally between two atoms, permitting each atom either by transfer or by sharing to be surrounded by a completed electron shell, usually of two or eight. Figure 1 summarizes the relations of Lewis's theory to earlier valence theories and to Bohr's theory of atomic structure, and points to its relations to the subsequent quantum-mechanical theories of molecular electronic structure. Lewis's theory showed electrons distributed in pairs shared by two atoms, and in unshared 'lone pairs'. As I have discussed elsewhere2, molecular orbital (MO) theory provides close counterparts to Lewis's shared pairs and