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The electron structure of atoms describes how electrons are distributed around the nucleus in discrete energy levels. These energy levels are called atomic orbitals and are characterized by quantum numbers that denote the principal energy, angular momentum, magnetic orientation, and spin of each electron. The arrangement of electrons in these orbitals determines an element’s chemical behaviour, including its valence, bonding patterns, and reactivity.

The basic rules governing electron placement are the Aufbau principle, Hund’s rule, and the Pauli exclusion principle. The Aufbau principle states that electrons occupy the lowest energy orbitals first, filling 1s, 2s, 2p, 3s, and so on. Hund’s rule dictates that within a degenerate set of orbitals (such as the three p orbitals), electrons populate separate orbitals singly before pairing up, maximizing the total spin. The Pauli exclusion principle prohibits two electrons in the same atom from having identical sets of quantum numbers, ensuring that each orbital can hold a maximum of two electrons with opposite spins.

Electron configurations are commonly written in shorthand notation, for example, helium is 1s² and neon is 1s² 2s² 2p⁶. Transition metals and heavier elements involve d and f orbitals and are represented by extended configurations such as [Ar] 3d⁶ 4s¹ for iron. Because the outermost electrons (valence electrons) are primarily responsible for chemical bonding, understanding electron structure is essential for predicting and explaining phenomena such as ionic versus covalent bonding, metallic conductivity, and magnetic properties.

The periodic table is organized largely by electron configurations; elements in the same column share similar valence shell arrangements, leading to analogous chemical properties. Advances in quantum chemistry and spectroscopy confirm the predictions of electron arrangement theory, enabling precise modelling of molecular behaviour and the design of novel materials and pharmaceuticals.