Introduction

Every molecule you will study in this chapter — from the simplest diatomic gas to a

complex ion — is held together by a chemical bond: the attractive force that binds atoms,

ions, or other constituents together within a chemical species. Free atoms almost never exist

on their own in nature (the noble gases are the one exception); instead, atoms combine into

groups with their own distinct, characteristic properties. Understanding why this combination

happens raises a cluster of natural questions: why do atoms bond at all? Why are only certain

combinations possible, while others never occur? Why do molecules end up with definite,

predictable shapes rather than a random arrangement of atoms?

This chapter works through four different theories, developed over time, that each answer these

questions from a different angle:

• The Kössel-Lewis approach — the earliest and simplest picture, built entirely around atoms achieving a stable, noble-gas-like electron count (the octet rule) by transferring or sharing electrons.
• The Valence Shell Electron Pair Repulsion (VSEPR) theory — extends the Lewis picture to predict a molecule's actual three-dimensional shape, based on how electron pairs around the central atom repel one another.
• Valence Bond (VB) theory — a quantum-mechanical treatment that explains bonding as the overlap of atomic orbitals, and introduces hybridisation to account for the directional properties of bonds.
• Molecular Orbital (MO) theory — a more complete quantum-mechanical picture that builds orbitals belonging to the whole molecule, rather than to any one atom, and correctly predicts properties (like the magnetism of $O_2$) that VB theory cannot.

These theories did not appear in isolation — each one grew directly out of the developing

understanding of atomic structure, electronic configuration, and the periodic table. The

throughline connecting all of them is a single physical idea: every system tends toward greater

stability, and forming a chemical bond is nature's way of lowering the energy of a collection of

atoms to reach that more stable state.

By the end of this chapter, you should be able to:

• explain the Kössel-Lewis approach to chemical bonding, including the octet rule and its limitations, and draw Lewis structures for simple molecules;
• describe the formation of ionic and covalent bonds;
• use VSEPR theory to predict the geometry of simple molecules;
• apply valence bond theory to explain covalent bond formation and the directional properties of bonds;
• describe the different kinds of hybridisation involving $s$, $p$, and $d$ orbitals, and draw the shapes of simple covalent molecules that result;
• describe molecular orbital theory as applied to homonuclear diatomic molecules; and
• explain the concept of the hydrogen bond.