Electron Configuration

Overview

Electron configuration describes how electrons are distributed among the orbitals of an atom. Understanding electron configurations helps predict chemical properties and reactivity.

Key Principles

1. Aufbau Principle

Electrons fill orbitals starting with the lowest energy level first.

2. Pauli Exclusion Principle

Each orbital can hold a maximum of 2 electrons with opposite spins.

3. Hund's Rule

Electrons occupy orbitals of the same subshell singly before pairing, all with the same spin.

Energy Level Order (Aufbau Diagram)

$$1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s < 4f < 5d < 6p < 7s < 5f < 6d < 7p$$

Memory Aid (Diagonal Rule)

1s
2s  2p
3s  3p  3d
4s  4p  4d  4f
5s  5p  5d  5f
6s  6p  6d
7s  7p

Follow the diagonals from top-right to bottom-left.

Notation

Spdf Notation

1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶

Noble Gas (Core) Notation

[Ar] 4s² 3d¹⁰ 4p⁶

The noble gas in brackets represents filled inner shells.

Orbital Diagram

1s: ↑↓   2s: ↑↓   2p: ↑↓ ↑↓ ↑↓

Subshell Capacities

Subshell	Orbitals	Max Electrons
s	1	2
p	3	6
d	5	10
f	7	14

Maximum electrons in subshell:

$$\text{Max electrons} = 2(2\ell + 1)$$

Examples

Oxygen (Z = 8)

1s² 2s² 2p⁴

Orbital diagram:

1s: ↑↓   2s: ↑↓   2p: ↑↓ ↑ ↑

Iron (Z = 26)

1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶

Or with noble gas notation:

[Ar] 4s² 3d⁶

Copper (Z = 29) - Exception

[Ar] 4s¹ 3d¹⁰  (not 4s² 3d⁹)

Half-filled and fully-filled d subshells are more stable.

Exceptions to Aufbau

Chromium (Z = 24)

Expected: [Ar] 4s² 3d⁴
Actual:   [Ar] 4s¹ 3d⁵

Copper (Z = 29)

Expected: [Ar] 4s² 3d⁹
Actual:   [Ar] 4s¹ 3d¹⁰

Half-filled (d⁵) and fully-filled (d¹⁰) subshells have extra stability.

Ion Electron Configurations

Cations (remove electrons from highest n first)

Fe:   [Ar] 4s² 3d⁶
Fe²⁺: [Ar] 3d⁶        (4s electrons removed first)
Fe³⁺: [Ar] 3d⁵

Anions (add electrons)

O:   1s² 2s² 2p⁴
O²⁻: 1s² 2s² 2p⁶  (same as Ne)

Valence Electrons

Electrons in the outermost shell (highest n value).

Element	Configuration	Valence e⁻
Na	[Ne] 3s¹	1
Mg	[Ne] 3s²	2
Al	[Ne] 3s² 3p¹	3
Cl	[Ne] 3s² 3p⁵	7

Paramagnetic vs Diamagnetic

• Paramagnetic: Has unpaired electrons (attracted to magnetic field)
• Diamagnetic: All electrons paired (slightly repelled by magnetic field)

O₂: ↑↓ ↑↓ ↑↓ ↑ ↑  →  Paramagnetic (2 unpaired)
Zn: [Ar] 4s² 3d¹⁰  →  Diamagnetic (all paired)