Bonding models: an introduction

A historical overview

  The foundations of modern chemical bonding theory were laid in 1916–1920 by G.N. Lewis and I. Langmuir who suggested that ionic species were formed by electron transfer, while electron sharing was important in covalent molecules.

In some cases, it was suggested that the shared electrons in a bond were provided by one of the atoms but that once the bond (sometimes called a coordinate bond) is formed, it is indistinguishable from a ‘normal’ covalent bond.  In a covalent species, electrons are shared between atoms.

In an ionic species, one or more electrons are transferred between atoms to form ions.  Modern views of atomic structure are, as we have seen, based largely on the applications of wave mechanics to atomic systems. Modern views of molecular structure are based on applying wave mechanics to molecules; such studies provide answers as to how and why atoms combine. The Schrodinger equation can be written to describe the behavior of electrons in molecules, but it can be solved only approximately. Two such methods are the valence bond approach, developed by Heitler and Pauling, and the molecular orbital approach associated with Hund and Mulliken:

•  Valence bond) VB) theory treats the formation of a molecule as arising from the bringing together of complete atoms which, when they interact, to a large extent retain their original character.
•  Molecular orbital (MO) theory allocates electrons to molecular orbitals formed by the overlap (interaction) of atomic orbitals.

Although familiarity with both VB and MO concepts is necessary, it is often the case that a given situation is more conveniently approached by using one or other of these models. We begin with the conceptually simple approach of Lewis for representing the bonding in covalent molecules.