Electron configuration

Reading the notation

In 1s² 2s² 2p⁶, the number gives the principal shell, the letter gives the subshell, and the superscript gives the number of electrons in that subshell. The capacities come from the number of orbitals: s has 1 orbital, p has 3, d has 5 and f has 7. Since each orbital can contain at most two electrons, the maximum populations are 2, 6, 10 and 14.

How electrons occupy orbitals

The ground-state configuration is the lowest-energy arrangement available to the atom. Three ideas work together: lower-energy orbitals are occupied first; no two electrons in one atom can have the same complete set of quantum numbers; and equal-energy orbitals are singly occupied before pairing begins.

For oxygen, 1s²2s²2p⁴ means that the three 2p orbitals do not fill as ↑↓, ↑↓, empty. The lower-energy arrangement is ↑↓, ↑, ↑: one pair and two unpaired electrons. That detail matters experimentally because unpaired electrons affect magnetism.

Noble-gas shorthand, ions and the 4s/3d issue

Sodium is 1s²2s²2p⁶3s¹, or [Ne]3s¹. Chlorine is [Ne]3s²3p⁵. The shorthand is not a different model; it simply replaces an unchanged inner core with the symbol of the preceding noble gas.

For main-group cations, electrons are usually removed from the highest occupied principal shell first. Na becomes Na⁺ by losing 3s¹. For transition metals, the commonly taught filling sequence can mislead: the 4s orbital is occupied before 3d in K and Ca, but after the 3d subshell becomes populated the relative orbital energies reorganize. Transition-metal cations therefore usually lose ns electrons before (n−1)d electrons.

Chromium and copper also show that the simple Aufbau diagram is an approximation: Cr is [Ar]3d⁵4s¹ and Cu is [Ar]3d¹⁰4s¹. The deeper reason is that 3d and 4s are close in energy, so electron–electron interactions can change the preferred ground state. “Half-filled shells are magically stable” is a mnemonic, not a complete explanation.

Chemical information encoded by configurations

Valence configurations organize the periodic table. Elements in the same main-group column have analogous outer configurations, which is why Li and Na both form +1 ions and F and Cl commonly form −1 ions. Blocks of the periodic table correspond to the type of subshell being filled.

Configurations also connect to atomic spectra, magnetism and bonding. An excited electron configuration differs from the ground state because an electron has been promoted to a higher-energy orbital. Spectroscopy can reveal these energy differences, while the number of unpaired electrons helps predict whether an isolated atom or ion is paramagnetic.

A configuration is therefore a model with experimental consequences, not just bookkeeping.

Photoelectron spectroscopy tests the configuration

Electron configurations can be checked experimentally with photoelectron spectroscopy (PES). Photons eject electrons from an atom or molecule, and the measured kinetic energies reveal how strongly the electrons were bound. Peaks at different binding energies correspond to different occupied shells and subshells.

For a simple atom such as neon, the 1s electrons appear at much larger binding energy than 2s or 2p because the 1s orbital lies close to the nucleus and experiences much less shielding. The relative peak areas also reflect how many electrons occupy each subshell. Thus the notation 1s²2s²2p⁶ predicts an experimental pattern rather than merely summarizing a rule.

PES also shows why “all electrons in one shell have the same energy” is false in many-electron atoms: 2s and 2p have different binding energies because their penetration and shielding differ.

Configuration, valence and chemical identity

Two species with the same electron count can have similar configurations but different chemistry because their nuclei differ. Ne, Na⁺ and F⁻ are all 10-electron species, yet their sizes and ionization energies differ because the same electron arrangement is held by different nuclear charges.

Conversely, elements in one group repeat analogous valence configurations even though their full configurations become longer. Li is [He]2s¹, Na is [Ne]3s¹ and K is [Ar]4s¹. That repeating ns¹ pattern is more chemically informative than memorizing every core electron.

Configurations are not complete many-electron states

A notation such as 2p² tells us which subshell is occupied, but it does not uniquely specify every possible way those electrons can couple their orbital and spin angular momenta. More advanced atomic spectroscopy distinguishes terms and levels within the same configuration. Hund’s rules then help order those levels.

For introductory chemistry, the configuration is the right level of description. It identifies the occupied orbitals and unpaired electrons; the finer structure becomes important when interpreting high-resolution spectra and transition-metal states.

Exercises

Sulfur

Write the ground-state configuration of sulfur, Z = 16, and count its outer-shell electrons.

Calcium ion

Write Ca²⁺ from neutral Ca = [Ar]4s².

Iron(III)

Neutral Fe is [Ar]3d⁶4s². Write Fe³⁺.