We define the location of every electron using three rules.
Aufbau Principle (Lowest Energy First)
Electrons fill the lowest energy subshell available first before occupying higher levels: 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p.
Pauli Exclusion Principle (Opposite Spins)
An orbital can hold a maximum of two electrons, and they must have opposite spins (↑↓).
Hund's Rule (Parallel Spins First)
Orbitals of equal energy (degenerate orbitals, like 2p) fill singly with parallel spins first before electrons pair up, minimising inter-electron repulsion.
THE EXCEPTIONS (Exam Favorite)
Half-filled (\(d^5\)) and fully-filled (\(d^{10}\)) sublevels offer extra stability.
[Ar] 4s² 3d⁴
[Ar] 4s¹ 3d⁵
[Ar] 4s² 3d⁹
[Ar] 4s¹ 3d¹⁰
Writing Configurations for Transition Metal Ions
Problem: Write the electron configuration for the Iron(III) ion (\(\text{Fe}^{3+}\)), where \(Z=26\).
1. Neutral Atom Configuration: \(\text{Fe} \; (26\text{e}^-) = 1\text{s}^2 2\text{s}^2 2\text{p}^6 3\text{s}^2 3\text{p}^6 4\text{s}^2 3\text{d}^6 = [\text{Ar}]\, 4\text{s}^2 3\text{d}^6\)
2. Ion Formation: Remove 3 electrons. Remove from 4s first (2 electrons), then 1 electron from 3d.
3. Final Result: \(\mathbf{\text{Fe}^{3+} = [\text{Ar}]\, 3\text{d}^5}\) (a stable half-filled d subshell).
Electron Loss in Transition Metals (4s before 3d)
When writing configurations for transition metal cations, always remove electrons from the 4s orbital BEFORE the 3d orbital. For example, neutral Iron (\(\text{Fe}\)) is \([\text{Ar}]\, 4\text{s}^2 3\text{d}^6\), but \(\text{Fe}^{2+}\) is \([\text{Ar}]\, 3\text{d}^6\) (the 4s electrons are lost first!).
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