An atom's electrons are arranged in energy levels, sublevels, and orbitals. The electron configuration describes that arrangement. It is written as a series of numbers and letters: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6, for example, is krypton. The number is the principal energy level. The letter is the sublevel type. The superscript is the number of electrons in that sublevel.

The rules for filling orbitals are straightforward. The Aufbau principle says electrons fill the lowest energy orbitals first. The Pauli exclusion principle says each orbital holds a maximum of two electrons with opposite spins. Hund's rule says that within a sublevel, electrons occupy separate orbitals with parallel spins before pairing up. These rules predict the electron configuration of most elements.

Electron configuration determines chemical properties. Elements in the same group have the same valence electron configuration, which is why they behave similarly. Sodium has one valence electron in the 3s orbital. Potassium has one in the 4s orbital. Both are alkali metals that react vigorously with water. The configuration also explains periodic trends in atomic radius, ionization energy, and electronegativity.

Exceptions to the rules occur in transition metals and some heavier elements. Chromium and copper, for example, have configurations that deviate from the predicted order because a half-filled or fully filled d sublevel is more stable. These exceptions are important for understanding the chemistry of transition metals.