Chemical Bonding

The concept of chemical bonding explains how individual atoms combine to form the vast array of substances we observe. Gilbert N. Lewis’s 1916 proposal of the covalent bond, where atoms share pairs of electrons to achieve a stable electron configuration (the ‘octet rule’), was a monumental step. This simple yet powerful idea, visualized with Lewis structures, provided the first clear picture of bonding in non-ionic compounds. Ionic bonds, in contrast, are formed by the complete transfer of one or more electrons from an atom with low electronegativity (typically a metal) to one with high electronegativity (a nonmetal), creating charged ions that are held together by electrostatic attraction.

Our understanding was deepened significantly by the application of quantum mechanics. Linus Pauling’s valence bond theory, introduced in the 1930s, described covalent bonds as the result of the overlap of atomic orbitals. This theory introduced the concepts of hybridization (e.g., sp, sp2, sp3) and resonance to explain observed molecular geometries and bond properties. An alternative and often more powerful approach is molecular orbital (MO) theory, which considers electrons to be delocalized across the entire molecule in molecular orbitals formed from the combination of atomic orbitals. MO theory is essential for explaining the properties of conjugated systems, magnetism, and excited states. Beyond these primary types, there are also metallic bonds, characterized by a ‘sea’ of delocalized electrons, and weaker intermolecular forces like hydrogen bonds and van der Waals forces, which are critical for the properties of liquids, solids, and biological macromolecules like DNA.