Electron Configuration

This lesson covers the arrangement of electrons in atoms, including shells, sub-shells, and orbitals, the rules governing electron filling, and the important exceptions at A-Level. Electron configuration underpins the whole of chemistry — from bonding to reactivity to periodicity.

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Shells, Sub-shells, and Orbitals

Shells (Principal Energy Levels)

Electrons are arranged in shells (principal energy levels), numbered n = 1, 2, 3, 4, etc. The shell number n determines:

• The maximum number of electrons in the shell: 2n²
• The number of sub-shells: n sub-shells

Shell (n)	Sub-shells	Maximum electrons (2n²)
1	1s	2
2	2s, 2p	8
3	3s, 3p, 3d	18
4	4s, 4p, 4d, 4f	32

Sub-shells

Each shell is divided into sub-shells labelled s, p, d, and f. These sub-shells have different energies (within the same shell, s < p < d < f).

Sub-shell	Number of orbitals	Maximum electrons
s	1	2
p	3	6
d	5	10
f	7	14

Orbitals

An orbital is a region of space around the nucleus where there is a high probability of finding an electron. Each orbital can hold a maximum of two electrons, which must have opposite spins (spin-up ↑ and spin-down ↓).

Shapes of orbitals:

• s orbitals: Spherical. Successive s orbitals (1s, 2s, 3s) are larger.
• p orbitals: Dumbbell-shaped (two lobes). There are three p orbitals per sub-shell, oriented along the x, y, and z axes (pₓ, p_y, p_z), at right angles to each other.
• d orbitals: More complex shapes — most have four lobes in a clover-leaf pattern. The dz² orbital has two lobes along the z-axis with a ring around the centre.

Key Point: All orbitals in the same sub-shell have the same energy — they are described as degenerate. For example, the three 2p orbitals are degenerate.

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Rules for Filling Orbitals

Three rules govern how electrons fill orbitals:

1. The Aufbau Principle

Electrons fill orbitals in order of increasing energy. The filling order is:

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

Note that 4s fills before 3d because 4s has a lower energy in neutral atoms. This is the most important exception to the simple numerical ordering.

2. Hund's Rule

Within a sub-shell, electrons occupy orbitals singly before pairing up. All singly occupied orbitals have electrons with the same spin (parallel spins). This minimises electron-electron repulsion.

For example, in nitrogen (7 electrons): 1s² 2s² 2p³ The three 2p electrons occupy the three 2p orbitals singly:

2p: [↑] [↑] [↑] — correct (Hund's rule) 2p: [↑↓] [↑] [ ] — incorrect (electrons should not pair before all orbitals have one)

3. Pauli Exclusion Principle

No two electrons in an atom can have the same set of four quantum numbers. In practice, this means each orbital can hold a maximum of two electrons, and they must have opposite spins.

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Writing Electron Configurations

Full Notation

The full electron configuration lists all sub-shells with the number of electrons as a superscript.

Examples:

Element	Z	Electron configuration
Hydrogen	1	1s¹
Helium	2	1s²
Lithium	3	1s² 2s¹
Carbon	6	1s² 2s² 2p²
Nitrogen	7	1s² 2s² 2p³
Oxygen	8	1s² 2s² 2p⁴
Neon	10	1s² 2s² 2p⁶
Sodium	11	1s² 2s² 2p⁶ 3s¹
Argon	18	1s² 2s² 2p⁶ 3s² 3p⁶
Potassium	19	1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹
Calcium	20	1s² 2s² 2p⁶ 3s² 3p⁶ 4s²
Scandium	21	1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹ 4s²
Titanium	22	1s² 2s² 2p⁶ 3s² 3p⁶ 3d² 4s²
Iron	26	1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁶ 4s²