Introduction

An atom consists of a tiny, dense, positively charged nucleus (containing protons and neutrons) surrounded by electrons occupying the much larger space around it. This picture, established by the scattering experiments and the discovery of the sub-atomic particles, tells us what an atom is made of, but it says almost nothing about how the electrons are arranged, why an atom absorbs and emits light only at certain sharp, characteristic wavelengths (its line spectrum), or why an electron does not simply spiral into the nucleus the way a classical charged particle orbiting another charge would be expected to.

Answering these questions required physics to move beyond classical mechanics and electromagnetism into quantum theory. This chapter builds that quantum mechanical picture of the atom in stages: Bohr's model, the first attempt to quantise the electron's orbit and the first model to correctly predict the hydrogen spectrum; the discovery that both light and matter show a dual, wave-and-particle character; Heisenberg's uncertainty principle, which shows why a fixed, classical orbit is not even in principle knowable; the Schrödinger wave equation, whose solutions -- the atomic orbitals -- replace the idea of a fixed orbit with a probability distribution; the quantum numbers that label every orbital; the characteristic shapes of the s, p and d orbitals; and, finally, the rules (Aufbau, Pauli, Hund) that fix how electrons actually fill those orbitals in the ground state of every atom, right up to the elements of the periodic table.