Electron ConfigurationJEE Main

Concept

Electrons fill orbitals from lowest energy upward — the Aufbau principle — obeying Pauli (max 2 per orbital, opposite spins) and Hund's rule (singly occupy degenerate orbitals first). The result is the atom's electron configuration.

Key formula

Derivation

The $(n+l)$ rule sets the energy order: 4s ($n+l=4$) fills before 3d ($n+l=5$). Filling Zn (Z=30) gives $1s^2\,2s^2\,2p^6\,3s^2\,3p^6\,4s^2\,3d^{10}$.

Chromium and copper break the pattern: a half-filled ($3d^5$) or fully-filled ($3d^{10}$) subshell is extra stable, so one 4s electron drops into 3d — giving $3d^5 4s^1$ and $3d^{10} 4s^1$.

Scenarios to explore

• Electron Configuration — Aufbau filling with Hund, Pauli and Cr/Cu anomalies.

Real-world applications

• Predicting valency and common oxidation states.
• Explaining periodic trends and magnetic properties.
• Rationalising colours of transition-metal complexes.

JEE exam tips

• Use the $(n+l)$ rule; for ties the lower $n$ fills first.
• When ionising transition metals, remove 4s electrons before 3d.
• Half-filled and fully-filled subshells confer extra stability.

Common mistakes

• Filling 3d before 4s (4s is lower in energy when filling).
• Pairing electrons before each degenerate orbital has one (violates Hund).
• Missing the Cr and Cu anomalies.

Exam traps to avoid

• Fe²⁺ is $3d^6$ (4s removed first), not $3d^4 4s^2$.
• The anomaly is about stability, not a violation of Aufbau energy ordering.