What is a bond?

Almost 100 years ago, the American chemist Gilbert Lewis proposed that chemical bonds arise from the sharing of pairs of electrons between atoms. In a 1916 paper,¹ he portrayed atoms as cubes with electrons at their corners, and argued that they accumulate an electron at every corner by sharing edges with other atoms. It was a dramatic shift from the prevailing notion of chemical bonding as an electrostatic interaction caused by the transfer of electrons between atoms. Not everyone liked the idea, but over the ensuing two decades Linus Pauling showed how electron sharing could be described by the new theory of quantum mechanics, making Lewis’ picture central to modern chemical bonding theory.

It still is. But since that time, variants on Lewis’ theme have proliferated. Dative bonds (now more generally called donor–acceptor bonds), where both electrons in the pair come from the same atom, and hydrogen bonds, where the interaction of a hydrogen atom with a lone pair is primarily electrostatic, are standard undergraduate fare. However, thanks partly to improved methods of determining molecular structure and partly to greater sophistication of quantum-chemical theory, new modes of bonding are still being discovered, and controversies continue to rage over whether or not a particular pair of atoms can be considered united by a bond. So is Lewis’ concept of a bond still fit for purpose in the 21st century?