Electron Configurations

Full Notes:

Orbitals and Electron Limits

Within subshells, electrons can only exist in certain regions of space. These regions of space are called ‘orbitals’.

Orbital key points:

• Each orbital is a region in space where there is a high chance of finding an electron.
• An orbital holds up to two electrons, and they must have opposite spins, shown as arrows pointing in opposite directions (↑↓).

Orbitals group into sublevels:

• s: 1 orbital → 2 electrons
• p: 3 orbitals → 6 electrons
• d: 5 orbitals → 10 electrons
• f: 7 orbitals → 14 electrons

Electron configurations are used to show which orbitals are occupied in a given atom (or ion) and how many electrons are in each.

Key Rules for Electron Configuration

Aufbau Principle
Electrons fill the lowest available energy orbitals first.
General filling order (up to Z = 36): 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p

Hund’s Rule
When electrons fill orbitals of equal energy (like the three p orbitals), they go into separate orbitals first with parallel spins. This reduces electron repulsion and increases stability.

Pauli Exclusion Principle
No two electrons in the same orbital can have the same spin. One must spin up (↑) and one must spin down (↓).

Arrow-in-Box Diagrams

Arrow-in-box diagrams can be used to show how orbitals are filled and relative energies. Each box = one orbital. Each arrow = one electron.

Example:Iron, Fe (atomic number of 26)

Writing Electron Configurations

When writing electron configurations, we use a standard notation:

The notation uses:

• A number for the shell (n)
• A letter for the sub-shell (s, p, d)
• A superscript for the number of electrons
• Inner (core) electron configurations can be represented using noble gas shorthand notation, i.e. [Ne] = 1s²2s²2p⁶

Examples:

• O (Z = 8) → 1s² 2s² 2p⁴
• Ca (Z = 20) → 1s² 2s² 2p⁶ 3s² 3p⁶ 4s²
• In shorthand: Ca = [Ar] 4s², Fe = [Ar] 3d⁶ 4s²

Energy and Electron Repulsion

Electron configurations are determined by:

• Energy levels (electrons fill lower energy sub-shells first)
• Electron–electron repulsion (electrons prefer to occupy separate orbitals in a sub-shell before pairing up, as this lowers repulsion)

This explains why 4s fills before 3d and also helps explain trends in ionisation energy.

Ions

For positive ions, remove electrons from the highest energy level first.

For transition metals, remove 4s electrons before 3d.

Example: Fe²⁺ (Z = 26):

• Neutral: [Ar] 4s² 3d⁶
• Fe²⁺: [Ar] 3d⁶

Exceptions to the Rules: Chromium and Copper

These are important to memorise:

• Chromium (Z = 24):
  Expected: [Ar] 4s² 3d⁴
  Actual: [Ar] 4s¹ 3d⁵
  (Half-filled d sublevel is more stable)
• Copper (Z = 29):
  Expected: [Ar] 4s² 3d⁹
  Actual: [Ar] 4s¹ 3d¹⁰
  (Filled d sublevel is more stable)

These exceptions occur due to the added stability of half-filled and fully-filled d orbitals.

Summary

• Each orbital holds up to 2 electrons with opposite spin.
• Electron filling follows Aufbau principle, Hund’s rule, and Pauli exclusion principle.
• Arrow-in-box diagrams show orbital occupation and electron spin.
• Use full and shorthand notation for configurations.
• Chromium and copper show exceptions due to stable half-filled and filled d sublevels.