In 1909, the New Zealand-born physicist Ernest Rutherford, working at Manchester, set up an experiment whose result he expected to be unremarkable. He fired a beam of alpha particles (helium nuclei) at a thin gold foil and measured how the particles deflected. The prevailing plum-pudding model of the atom (J. J. Thomson, 1904) said the atom was a diffuse positive sphere with electrons embedded like raisins; alpha particles should pass through with minimal deflection. Most did. But one in eight thousand bounced almost straight back. Rutherford recalled the moment as the most incredible event of my life. "It was as though you had fired a fifteen-inch shell at a piece of tissue paper and it came back and hit you." The conclusion was inescapable: the atom has a tiny dense nucleus, where almost all the mass and positive charge is concentrated, with the electrons orbiting at large distances. The nucleus, Rutherford calculated, is some hundred thousand times smaller than the atom. Atoms are mostly empty space.
An atom consists of a tiny dense nucleus (made of protons and neutrons, called nucleons) of size roughly 10⁻¹⁵ m, surrounded by an electron cloud of size roughly 10⁻¹⁰ m. Almost all the mass is in the nucleus (a proton is ~1836 times heavier than an electron); almost all the size is in the electron cloud. The nucleons are bound by the strong nuclear force (binds protons against their mutual electromagnetic repulsion); the electrons are bound by the electromagnetic attraction of the nucleus. The number of protons (the atomic number Z) defines the chemical element. The number of neutrons can vary, giving isotopes of the same element. Bohr's model (1913) — a transitional theory before full quantum mechanics — had electrons in quantized circular orbits with discrete allowed radii and energies, explaining the line spectrum of hydrogen. The full quantum-mechanical picture, built from solutions to the Schrödinger equation in a Coulomb potential, replaces orbits with orbitals — three-dimensional probability densities for the electron's position. Each orbital is characterized by quantum numbers: the principal n (energy level: 1, 2, 3, ...), the angular momentum l (shape: s, p, d, f, ...), the magnetic m_l (orientation), and the spin m_s (intrinsic angular momentum: ±½). The Pauli exclusion principle (1925) — no two electrons can occupy the same quantum state — forces electrons into successive shells of increasing energy. The first shell holds 2 electrons (1s); the second holds 8 (2s, 2p); the third holds 18 (3s, 3p, 3d); and so on. This shell structure is what gives the periodic table its rows and columns. Hund's rule, the Aufbau principle, Madelung's rule — the empirical patterns of which orbitals fill in what order — are direct consequences. Spectroscopy (the study of the wavelengths atoms emit and absorb) is a precision probe of orbital structure and the foundation of analytical chemistry, astrophysics, and atomic clocks.
Atomic structure is the bridge between physics and chemistry. The periodic law — already in our system — is a direct consequence of atomic shell structure. Materials science, semiconductor design, drug design, biochemistry — all rest on atomic-scale understanding. Atomic clocks (cesium clocks define the SI second; modern optical lattice clocks are roughly 10¹⁸ times more precise) exploit ultra-stable atomic transitions. Mass spectrometry identifies molecules by their isotopic composition; X-ray crystallography and electron microscopy image atomic arrangements directly. Nuclear medicine uses radioactive isotopes for diagnosis (PET, SPECT) and treatment (radiotherapy). Nuclear power and weapons exploit nuclear binding energies through fission and fusion. The little nucleus Rutherford glimpsed in 1909 is the centre of a tower of subsequent science.