Electron Configuration

What is it?

An electron configuration shows how an atom’s electrons are arranged in shells and subshells. For example, oxygen (8 electrons) is:

1s^2\, 2s^2\, 2p^4

Read it as: 2 electrons in the 1s subshell, 2 in 2s, and 4 in 2p.

The number is the shell (energy level), $n$ = 1, 2, 3…
The letter is the subshell type: s, p, d or f.
The superscript is the number of electrons in that subshell.

Key idea

Electrons fill the lowest-energy subshells first. Each subshell holds a fixed maximum: s = 2, p = 6, d = 10, f = 14 electrons.

Why does it matter?

It explains the periodic table. Elements in the same group have similar outer (valence) configurations, which is why they behave alike.
It predicts chemistry. The number of valence electrons determines how an atom bonds and what ions it forms.
It explains colour, magnetism and more. Unpaired electrons and partly filled d subshells give many transition-metal compounds their colours and magnetic properties.

How does it work?

1. Shells, subshells and orbitals

Subshell	Orbitals	Maximum electrons
s	1	2
p	3	6
d	5	10
f	7	14

Each orbital holds at most 2 electrons with opposite spins (the Pauli exclusion principle).

2. The filling order (Aufbau principle)

Subshells fill in order of increasing energy:

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

Notice that 4s fills before 3d. The diagonal diagram below is an easy way to remember the order.

3. Hund’s rule

Within a subshell, electrons occupy separate orbitals with parallel spins before any orbital gets a second electron. For nitrogen’s three 2p electrons: [↑] [↑] [↑], not [↑↓] [↑] [ ].

4. Noble-gas shorthand

Replace the inner (core) electrons with the previous noble gas in square brackets:

Sodium: $1s^2\, 2s^2\, 2p^6\, 3s^1$ = [Ne] 3s¹
Iron: [Ar] 3d⁶ 4s² (also written [Ar] 4s² 3d⁶)

5. Ions

Anions: add electrons. $\ce{Cl-}$ : [Ne] 3s² 3p⁶ = [Ar].
Main-group cations: remove electrons from the outermost subshell. $\ce{Na+}$ : [Ne].
Transition-metal cations: remove 4s electrons first, then 3d. $\ce{Fe^2+}$ : [Ar] 3d⁶; $\ce{Fe^3+}$ : [Ar] 3d⁵.

Think of it like this

Think of a hotel filling rooms from the ground floor up. Guests (electrons) take the cheapest rooms (lowest energy) first. Hund’s rule is like passengers on a bus: people take empty double seats on their own before anyone sits next to a stranger.

More precisely

A few elements break the simple filling order because half-filled and completely filled d subshells are especially stable: chromium is [Ar] 3d⁵ 4s¹ (not 3d⁴ 4s²) and copper is [Ar] 3d¹⁰ 4s¹ (not 3d⁹ 4s²). Also, once the 3d subshell contains electrons, its energy drops below 4s. That’s why transition metals lose their 4s electrons first when they form ions.

Visualise it

Diagonal rule diagram for the order of filling subshells. Subshells are arranged in rows by shell: 1s; 2s 2p; 3s 3p 3d; 4s 4p 4d 4f; 5s 5p 5d 5f; 6s 6p 6d; 7s 7p. Diagonal arrows running from upper right to lower left give the filling order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p. — Follow the diagonal arrows to get the filling order.

Worked example

Worked example: Writing a configuration

Question: Write the full and noble-gas configurations of sulfur (Z = 16).

16 electrons, filled in order: 1s² (2), 2s² (4), 2p⁶ (10), 3s² (12), 3p⁴ (16)
Full: $1s^2\, 2s^2\, 2p^6\, 3s^2\, 3p^4$
Noble-gas form: [Ne] 3s² 3p⁴: 6 valence electrons, two of them unpaired (Hund’s rule).

Worked example: A transition-metal ion

Question: Write the configuration of $\ce{Fe^3+}$ (Fe: Z = 26).

Fe atom: [Ar] 3d⁶ 4s²
Remove 3 electrons, 4s first: remove both 4s electrons, then one 3d electron.
$\ce{Fe^3+}$ : [Ar] 3d⁵ (five unpaired electrons, one in each 3d orbital)

Common mistake

Common mistake: Filling 3d before 4s

Potassium is [Ar] 4s¹, not [Ar] 3d¹. The 4s subshell fills before 3d.

Common mistake: Removing 3d electrons first from transition metals

When transition metals form ions, the 4s electrons are removed first. $\ce{Fe^2+}$ is [Ar] 3d⁶, not [Ar] 3d⁴ 4s².

Common mistake: Pairing electrons too early

Hund’s rule: put one electron in each orbital of a subshell before pairing. Oxygen’s 2p⁴ is [↑↓] [↑] [↑], with two unpaired electrons.

Notation note

Superscripts are sometimes written on the line, e.g. 1s2 2s2 2p4, when superscripts aren’t available.
The order of writing 3d and 4s varies between textbooks ([Ar] 3d⁶ 4s² or [Ar] 4s² 3d⁶). Both are accepted.
“Aufbau” is German for “building up”.

Remember this

Subshell capacities: s 2, p 6, d 10, f 14; each orbital holds 2 electrons.
Fill in order (4s before 3d); Hund’s rule: spread out before pairing.
Noble-gas shorthand: [previous noble gas] + outer electrons.
Transition-metal ions lose 4s electrons first. Exceptions: Cr and Cu.

Practise this topic with flashcards and a quiz at chemistryclarity.com/chemistry/electron-configuration/