Writing electron configurations

The big idea: An electron configuration is a map of how an atom's electrons are spread across its sub-shells (1s, 2s, 2p, 3s, …).

Electrons fill those sub-shells in a fixed way, set by three rules. Get the rules right and you can write the configuration of any atom or ion in the first few rows of the periodic table.

- Aufbau — fill the lowest energy first. - Pauli — at most 2 electrons per orbital, with opposite spins. - Hund — fill orbitals singly first, then pair.

The vocabulary: - Sub-shell — a set of orbitals of the same type and energy: s (1 orbital), p (3 orbitals), d (5 orbitals). - Orbital — a region holding up to 2 electrons. Drawn as a box. - Ground state — the lowest-energy arrangement, the one these rules give. Anything else is an excited state. - The superscript (e.g. the 4 in 2p⁴) is the number of electrons in that sub-shell.

Drawing electrons in boxes makes the three rules visible. Each box is an orbital; an arrow is an electron; the direction of the arrow is its spin. Fill from the bottom up (Aufbau).

The filling order to memorise: The sub-shell energy order across the first four rows is:

1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p

Note the swap: 4s fills before 3d. The maximum electrons are 2 (s), 6 (p) and 10 (d), matching the number of boxes × 2.

Why nitrogen looks like that: Nitrogen has 7 electrons: 1s² 2s² 2p³. The three 2p electrons could crowd into one or two boxes — but Hund's rule says spread them out singly with parallel spins first. So the 2p sub-shell is ↑ ↑ ↑, not ↑↓ ↑.

Then pairing begins: Oxygen has one more electron than nitrogen (1s² 2s² 2p⁴). The 2p sub-shell is full of singly-filled boxes, so the extra electron must now pair up (Pauli: opposite spin) — giving ↑↓ ↑ ↑.

A full configuration lists every sub-shell. A condensed (or core) configuration replaces the inner electrons with the previous noble gas in square brackets, then lists only the outer electrons — quicker and clearer.

Ions — add or remove from the outermost shell: To write an ion configuration, start from the atom, then:

- Negative ion (gained electrons) → add electrons to the next available sub-shell. e.g. O → O²⁻ adds 2 to give 1s² 2s² 2p⁶. - Positive ion (lost electrons) → remove electrons from the highest occupied main shell (largest n) first.

For transition metals this matters: the 4s electrons leave before the 3d electrons, even though 4s filled first.

The 4s-out-first trap: Iron is [Ar] 3d⁶ 4s². To make Fe²⁺, remove the two 4s electrons (highest n), not two 3d — giving [Ar] 3d⁶. Removing 3d by mistake is the classic lost mark.

The two exceptions: chromium and copper: A half-full or full d sub-shell is extra stable, so two elements break the simple Aufbau order:

- Chromium is [Ar] 3d⁵ 4s¹ (not 3d⁴ 4s²) — one electron promotes to give a half-full 3d⁵. - Copper is [Ar] 3d¹⁰ 4s¹ (not 3d⁹ 4s²) — to give a full 3d¹⁰.

These two are worth memorising; the IB expects them.

How this is tested: S1.3.3 shows up two ways.

- Paper 1A (MCQ): pick the correct orbital diagram for an atom (a Hund check), or the right condensed configuration — copper and chromium are favourite exceptions to spot. - Paper 2: a state or deduce question — write the full configuration of an atom, or deduce the configuration of an ion.

The deduce-an-ion question is the one that separates marks: you must remove electrons from the highest main shell first (4s before 3d).

Scoring the deduce-the-ion mark: Write the neutral atom first, then strike out the electrons you remove. For a transition-metal cation, take the 4s electrons out before any 3d — that is the marking point examiners look for.