Explain why the first ionisation energy of oxygen is lower than that of nitrogen, even though oxygen has a higher nuclear charge.

IB DP Chemistry Standard Level (2023 syllabus) — S1.3 Electron configurations · Explain · 4 marks · View as Markdown

Written & reviewed by James Millett — Biology (Imperial College London), PGCE Science (University of Cambridge).

Model answer (4 marks)

1. Nitrogen’s 2p sub‑level is half‑filled (2p^3) – one electron in each orbital, giving a highly stable configuration.
2. This stability means more energy is needed to remove an electron from N, raising its first ionisation energy.
3. Oxygen has 2p^4; the fourth electron must pair in an already occupied orbital.
4. The resulting electron–electron repulsion makes it easier to remove an electron, so O’s first ionisation energy is lower than N’s.

Examiner tips

Use the term ‘half‑filled sub‑level’ and ‘electron–electron repulsion’ – these are key phrases the mark scheme looks for.
Show the comparison clearly: stability of N → higher IE, repulsion in O → lower IE.
Keep the answer concise – 4 points, no extra explanation.

Common mistakes

Confusing ‘higher nuclear charge’ with ‘higher ionisation energy’ – the question asks why IE is lower despite higher charge.
Failing to mention the half‑filled 2p configuration of nitrogen.
Not explaining that the paired electron in oxygen increases repulsion and lowers IE.

Mark scheme (4 marks)

Nitrogen has a half-filled 2p sub-level (one electron in each of the three 2p orbitals), which is a particularly stable arrangement.
This half-filled arrangement gives nitrogen extra stability, so more energy is required to remove an electron from nitrogen than expected.
In oxygen, the fourth 2p electron must pair up in an already occupied orbital, causing electron–electron repulsion.
This repulsion between paired electrons in oxygen makes it easier to remove one of them, lowering oxygen's first ionisation energy relative to nitrogen.

Key terms in this question

first ionisation energy