Writing Electron Configurations

The three rules with a worked example

An electron configuration is the address list of an atom's electrons: which orbitals they occupy, in energy order. Oxygen (Z = 8): $1s^2 2s^2 2p^4$. Read it as: two electrons in 1s, two in 2s, four in 2p.

Three rules build every configuration:

1. Aufbau principle — fill lowest-energy orbitals first: 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p…
2. Pauli exclusion principle — each orbital holds at most two electrons, with opposite spins.
3. Hund's rule — in a subshell, place one electron per orbital (parallel spins) before pairing any.

Apply to nitrogen (Z = 7): 1s² (Aufbau, Pauli), 2s², then 2p³ with one electron in each p orbital — $1s^2 2s^2 2p_x^1 2p_y^1 2p_z^1$. Hund's rule is why nitrogen has three unpaired electrons, which is also why it forms three bonds.

Noble-gas shorthand

Full configurations get long, so chemists abbreviate with the preceding noble gas. Chlorine (Z = 17): full $1s^2 2s^2 2p^6 3s^2 3p^5$ becomes $[Ne]3s^2 3p^5$ — neon's configuration ($1s^2 2s^2 2p^6$) is packed into [Ne], and only the valence part is written out.

The shorthand is honest: the bracketed core really is chemically inert, identical to the noble gas. What remains outside the bracket is exactly the valence configuration that determines the element's chemistry. Reading $[Ne]3s^2 3p^5$ you instantly see: third-period, seven valence electrons, halogen behavior.

Valence configurations compress further: main-group elements are described by $ns^x np^y$ alone. This is the language periodic trends are written in.

The famous exceptions: chromium and copper

Two elements famously break the simple order. Chromium (Z = 24) is $[Ar]4s^1 3d^5$, not $[Ar]4s^2 3d^4$; copper (Z = 29) is $[Ar]4s^1 3d^{10}$, not $[Ar]4s^2 3d^9$. In each case one 4s electron promotes into 3d to achieve a half-filled ($d^5$) or filled ($d^{10}$) subshell.

Why? Half-filled and filled subshells gain extra stability from symmetric electron distribution and exchange energy — enough to outweigh the small 4s/3d energy gap. The exceptions prove the deeper principle: configurations minimize total energy, and the filling order is a usually-reliable guide to that minimum, not a law.

For exams, memorize Cr and Cu. For understanding, remember the principle: whenever subshell energies are close, stability subtleties can reshuffle the last electrons.

Configurations of ions

Ions reconfigure by gaining or losing electrons, and the removal order surprises many students: remove from the highest principal quantum number first, regardless of filling order. Iron is $[Ar]4s^2 3d^6$; Fe²⁺ loses the 4s electrons first, giving $[Ar]3d^6$ — not $[Ar]4s^2 3d^4$.

The reason: once occupied, the 3d orbitals drop below 4s in energy (greater nuclear attraction at smaller average radius), so 4s becomes the outermost, highest-energy subshell — the first to empty. Fe³⁺ goes one further: $[Ar]3d^5$, a half-filled d subshell, which is why the +3 state is so stable for iron.

Anions simply add electrons to the next available orbital: oxygen $[He]2s^2 2p^4$ gains two to become O²⁻ = $[He]2s^2 2p^6$, neon's configuration. Main-group ions almost always land on noble-gas configurations — the octet rule restated in orbital language.

Putting it together: identify and predict

An unknown element has the valence configuration $4s^2 4p^4$. Identify it and predict its ion. The 4p⁴ valence with filled 4s places it in period 4, group 16: selenium. As a nonmetal one group from the halogens, it gains two electrons to reach krypton's configuration, forming Se²⁻.

Check: selenium Z = 34, configuration $[Ar]4s^2 3d^{10} 4p^4$ — 34 electrons ✓. Se²⁻ adds two: $[Ar]4s^2 3d^{10} 4p^6$ = krypton core ✓. This is configuration thinking in action: from six symbols of shorthand to element identity, ion charge, and chemical family.

Common mistakes

Real-world applications

• Predicting ion charges from configurations (which electrons are lost first)

• Explaining periodic trends via shielding and effective nuclear charge

• Identifying elements from spectroscopic or configuration data

• Understanding magnetism: unpaired electrons make atoms paramagnetic