Introduction

“Covalency” refers to the partial sharing of electrons that leads to mutual attraction of atoms, the profound and characteristic mystery of chemical bonding. The concept of covalency largely evolved from classical organic chemistry (the study of compounds of carbon, hydrogen, and nearby main-group elements) and achieved considerable maturity in the mid nineteenth century, long before Rutherford's model of atomic structure made it possible to consider the deeper electronic implications of this concept. The latter step was achieved most notably by G. N. Lewis, who showed how the octet rule and shared-electron-pair concepts could provide a comprehensive rationalization of structural bonding principles. The elementary Lewis-structure picture of nonbonding (one-center) and bonding (two-center) valence electron pairs forms the starting point for practically all aspects of molecular bonding in s- and p-block elements, the subject of the present chapter. Extension of covalency concepts to d-block elements will be considered in Chapter 4.

Let us briefly outline the topics to be addressed, which span a rather wide range of covalent and noncovalent “effects” in the bonding of s/p-block elements. The localized Lewis-structure picture of covalent and polar covalent bonding is described in Section 3.2, starting from the simplest diatomic species (e.g., H2 +, H2, and dialkali analogs) and proceeding through singly and multiply bonded polyatomic species. Inadequacies of the single-Lewis-structure picture (“delocalization effects”) are described in Section 3.3, including conjugative and aromatic “resonance” effects, which are conventionally represented as contributions from additional Lewis structures. Weaker “hyperconjugative” delocalizations are described in Section 3.4, including those responsible for internal rotation barriers, anomeric effects, and other stereoelectronic phenomena.