Electron Configuration

Full Notes

Electron configurations and orbital shapes are covered in more detail at Electron Orbitals and Electron Configurations.

This page is just what you need to know for AQA A-level Chemistry :)

Electrons in Atoms

Electrons in an atom are arranged in energy levels (shells) labelled by the principal quantum number n (n = 1, 2, 3…). Each shell contains subshells (s, p, d; and f from n = 4). Each subshell consists of orbitals — regions of space with high probability of finding an electron — and each orbital holds a maximum of 2 electrons with opposite spins.

• s subshell → 1 orbital → holds 2 electrons
• p subshell → 3 orbitals → holds 6 electrons
• d subshell → 5 orbitals → holds 10 electrons

Electrons fill from the lowest energy level upwards (Aufbau principle).

Order of filling: 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p

Electron Configurations (Full and Shorthand)

Electron configurations show how electrons are arranged within atoms or ions.

The notation uses: a number for the shell (n), a letter for the subshell (s, p, d), and a superscript for the number of electrons.

Examples Full configuration:
O (Z = 8) → 1s² 2s² 2p⁴
Ca (Z = 20) → 1s² 2s² 2p⁶ 3s² 3p⁶ 4s²

Shorthand notation uses the nearest noble gas to represent inner (core) electrons:
Ca = [Ar] 4s²
Fe = [Ar] 3d⁶ 4s²

Matt’s exam tip

When writing electron configurations, double‑check that you’ve filled subshells in the correct order and that your total number of electrons matches the element or ion’s charge. For ions, remove electrons from the outermost shell (highest principal energy level) - see below.

Electron Configuration of Ions

When atoms form ions:

• Cations → lose electrons, starting with the highest energy level.
• Anions → gain electrons into the next available orbital.

Examples:

• Na (Z=11): 1s² 2s² 2p⁶ 3s¹ → Na⁺ = 1s² 2s² 2p⁶
• Cl (Z=17): 1s² 2s² 2p⁶ 3s² 3p⁵ → Cl⁻ = 1s² 2s² 2p⁶ 3s² 3p⁶

First Ionisation Energy

First Ionisation Energy has been covered in more detail here.

Definition: Energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1⁺ ions.

Equation: X(g) → X⁺(g) + e⁻

Successive ionisation energies involve further electrons being removed:

• 2nd IE: X⁺(g) → X²⁺(g) + e⁻
• 3rd IE: X²⁺(g) → X³⁺(g) + e⁻

Trends in Ionisation Energies

Period 3 (Na–Ar): Ionisation energy increases across the period.
Reason: more protons in the nucleus means stronger nuclear attraction to outermost electrons whilst shielding stays similar.

Exceptions:

• Al lower than Mg – outer electron in 3p orbital (higher energy).
• S lower than P – electron pair repulsion in 3p orbital.

Group 2 (Be–Ba): Ionisation energy decreases down the group.

Reason: atomic radius increases and more shielding means weaker attraction between nucleus and outermost electron.

Successive Ionisation Energy Graphs

Sharp increases in successive ionisation energies give evidence of shell structure.

For example, a large jump between 1st and 2nd ionisation energies means the second electron is in an inner shell.

Summary

• Electron Configuration: Arrangement of electrons in shells and sub-shells. 1s²2s²2p⁶...etc
• Ionisation Energy: Energy needed to remove an electron from a gaseous atom.
• First Ionisation Energy: X(g) → X⁺(g) + e⁻
• Successive Ionisation Energies: Removal of multiple electrons, one at a time.
• Trend Across Period 3: Increases overall, with drops at Al and S.
• Trend Down Group 2: Decreases due to increasing atomic radius and shielding.