Electronic configurations

The electronic configuration shows how the electrons of an atom are arranged in its orbitals. Three rules decide it:

• Aufbau principle ('building up'): in the ground state, electrons fill orbitals in order of increasing energy: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p …
• Pauli exclusion principle: no two electrons in an atom can have the same four quantum numbers. So an orbital holds at most two electrons, with opposite spins.
• Hund's rule of maximum multiplicity: within a subshell, electrons go one into each orbital, with parallel spins, before any pairing starts. Like passengers on a bus who each take an empty seat before anyone sits next to a stranger.

Examples: oxygen (8 electrons) is 1s² 2s² 2p⁴. Carbon (6 electrons) is 1s² 2s² 2p²; by Hund's rule, its two 2p electrons sit in separate orbitals, so carbon has two unpaired electrons.

Long configurations can be shortened by writing the previous noble gas in brackets: sodium is [Ne] 3s¹, and potassium is [Ar] 4s¹.

Atoms and ions with the same number of electrons are isoelectronic: Na⁺, Mg²⁺ and Ne all have 10 electrons, with the configuration 1s² 2s² 2p⁶.

Half-filled (p³, d⁵, f⁷) and completely filled (p⁶, d¹⁰, f¹⁴) subshells are extra stable, because of their symmetry and their larger exchange energy (electrons with parallel spins can swap places, which lowers the energy). This explains some exceptions to the Aufbau order. Copper, for example, is [Ar] 3d¹⁰ 4s¹, not 3d⁹ 4s²: moving one electron from 4s to 3d gives a completely filled d subshell. Chromium is another such exception, gaining a half-filled d subshell in the same way.