Electron configuration

Electron configuration records which atomic orbitals contain an atom’s electrons and how those electrons are distributed among them. ¹ The notation begins with a principal shell number, followed by a subshell letter: s, p, d, or f. A superscript gives the number of electrons in that subshell, so 1s² 2s² 2p⁶ represents ten electrons and is the ground-state configuration of neon. The subshell letters correspond to the orbital angular-momentum quantum number: s has ℓ = 0, p has ℓ = 1, d has ℓ = 2, and f has ℓ = 3.

Each individual orbital is described by a magnetic quantum number and can contain no more than two electrons, whose spins must be opposite. Thus an s subshell contains one orbital and holds two electrons, p contains three orbitals and holds six, d contains five and holds ten, and f contains seven and holds fourteen. ² Configurations may be written in full or abbreviated with a preceding noble-gas core, such as [Ne] 3s² 3p⁵ for chlorine.

Ground-state configurations are constructed by combining the Aufbau principle, the Pauli exclusion principle, and Hund’s rule. The Aufbau principle places electrons in orbitals of progressively higher energy, commonly represented by the order 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, and onward. ³ The Pauli principle prohibits two electrons in one atom from having the same complete set of four quantum numbers; an orbital therefore holds at most two electrons with opposite spin. Hund’s rule places electrons singly in degenerate orbitals before pairing, producing parallel spins when possible.

The familiar filling sequence is a practical guide rather than a universal law that fixes orbital energies in every circumstance. Orbital energies depend on electron–electron repulsion and the atom’s charge, and the order changes for ions and highly excited states. Electron configurations consequently describe an energy state, not merely a mechanical arrangement of particles. Spectroscopic measurements and quantum calculations provide the more precise account of those states. ¹

Electron configuration explains why the periodic table is divided into s, p, d, and f blocks. Elements in a group often share the same outer-shell pattern, so their valence electrons give a first approximation to similar chemical behavior. Alkali metals commonly end in ns¹, halogens in ns²np⁵, and noble gases in filled valence shells. ⁴ When atoms form ions, electrons are removed or added to produce a new configuration; for transition-metal cations, electrons are generally removed from the ns orbital before the (n−1)d orbital, even though ns filled first in the neutral atom.

Some configurations depart from the simplest Aufbau prediction because subshell energies are close and electron correlation stabilizes particular arrangements. Chromium is conventionally written [Ar] 3d⁵ 4s¹ rather than [Ar] 3d⁴ 4s², while copper is [Ar] 3d¹⁰ 4s¹ rather than [Ar] 3d⁹ 4s². These patterns should be treated as measured or calculated ground-state results, not as a general rule that half-filled or filled subshells always prevail. ³

Electron configuration also applies to excited states, atoms with unusual charge, and species whose electrons are not well represented by a single simple configuration. An excited atom can have an electron promoted to a higher orbital, creating a configuration that differs from its ground state; the resulting transitions produce characteristic spectral lines. The electron configuration of an open-shell atom may be expressed through several coupled terms, because electron–electron interactions split states with the same broad orbital occupancy into different energies. ⁵

For heavy atoms, relativistic effects and spin–orbit coupling become increasingly significant, so textbook orbital labels remain useful shorthand but do not capture every detail of the wavefunction. In spectroscopy, databases such as the NIST Atomic Spectra Database list observed and calculated energy levels, configurations, and transitions. ⁵ In molecules, electrons occupy molecular orbitals formed from atomic orbitals; atomic configurations remain the starting point, but molecular symmetry and bonding determine the final distribution. The notation is therefore both a concise chemical language and an approximation whose precision depends on the question being asked.

Atomic orbital

A one-electron spatial wavefunction, or a region described by it, characterized by quantum numbers and capable of holding two electrons with opposite spins.

Subshell

A group of orbitals in one principal shell sharing the same angular-momentum quantum number, designated s, p, d, or f.

Valence electron

An electron in an outer or chemically active shell that can participate in bonding and reactions.

Aufbau principle

The practical rule that ground-state electrons occupy available orbitals in order of increasing energy.

Hund’s rule

The rule that electrons occupy degenerate orbitals singly, with parallel spins, before pairing.

Noble-gas notation

An abbreviated configuration that replaces an inner closed-shell core with the symbol of the preceding noble gas.