OCR A Level Chemistry A: Electron structure (H432) – Revision Notes

These revision notes cover 2.2.1 Electron structure from Module 2 (Foundations in chemistry) of the OCR A Level Chemistry A (H432) specification, Version 3.1 (May 2026), for first assessment in 2017. They condense every statement in 2.2.1(a) to (d). H432 is the full A level with no tiers. For full explanations and worked examples, read the electron structure study guide first.

Once these notes feel secure, test yourself with the electron structure practice questions. The course hub and course checklist show the rest of H432, atomic structure and isotopes covers the content just before this, and a free diagnostic will show up weak spots quickly.

Key definitions

• Shell (energy level): a main energy level, numbered n = 1, 2, 3, 4 outwards from the nucleus.
• Sub-shell: a division of a shell, labelled s, p or d (the fourth shell also has f).
• Atomic orbital (2.2.1(b)(i)): a region around the nucleus that can hold up to two electrons, with opposite spins.
• Electron configuration: the arrangement of electrons in sub-shells, written in sub-shell notation such as 1s² 2s² 2p⁴ for oxygen.
• Isoelectronic: having the same electron configuration (for example Mg²⁺, Ne and F⁻).

Numbers to know: 2.2.1(a) and (b)(iii)

Shells

Shell	Sub-shells	Orbitals	Maximum electrons
n = 1	1s	1	2
n = 2	2s, 2p	4	8
n = 3	3s, 3p, 3d	9	18
n = 4	4s, 4p, 4d, 4f	16	32

Rule of thumb: shell n holds 2n² electrons in n² orbitals.

Sub-shells

Sub-shell type	Orbitals in it	Electron capacity
s	1	2
p	3	6
d	5	10

Orbital shapes: 2.2.1(b)(ii)

• s: a sphere centred on the nucleus. Higher shells give bigger spheres.
• p: a dumbbell, two lobes either side of the nucleus. Three p-orbitals per sub-shell, arranged at 90° to each other along the x, y and z axes.

Filling rules: 2.2.1(c)

1. Lowest energy first. Order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p.
2. Two per orbital, opposite spins. Draw them as ↑↓.
3. Singly before pairing. In a p- or d-sub-shell, put one electron in each orbital (same spin direction) before any orbital gets a second.

Memory hook for the tricky step: after 3p comes 4s, then 3d, then 4p.

Electrons in boxes: titanium, Z = 22

Configuration: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d² 4s²

1s [↑↓]   2s [↑↓]   2p [↑↓][↑↓][↑↓]
3s [↑↓]   3p [↑↓][↑↓][↑↓]
4s [↑↓]   3d [↑ ][↑ ][  ][  ][  ]

The two 3d electrons sit in separate orbitals with parallel spins, leaving three d-orbitals empty. Titanium has two unpaired electrons.

Nitrogen and oxygen side by side

Atom	Configuration	2p boxes	Unpaired electrons
N, Z = 7	1s² 2s² 2p³	[↑ ][↑ ][↑ ]	3
O, Z = 8	1s² 2s² 2p⁴	[↑↓][↑ ][↑ ]	2

Oxygen’s fourth 2p electron has to pair, because all three 2p orbitals already hold one electron. That pairing matters again in 3.1.1, where it helps explain ionisation energy patterns across Period 2.

Why the model was refined (HSW1, HSW7)

The additional guidance links 2.2.1 to how scientific models develop. The simple shell model (2, 8, 8) cannot explain why potassium starts the fourth shell before the third is full. Splitting shells into sub-shells and orbitals with different energies explains this, and also explains patterns in ionisation energies. If a question asks why the model changed, the answer is that new evidence did not fit the older model, so the model was refined to account for it.

Method in steps: configuration of an atom, 2.2.1(d)(i)

1. Find Z. That is the number of electrons.
2. Fill sub-shells in energy order, keeping a running total.
3. Stop when the running total equals Z. The last sub-shell may be part-full.
4. Write it out with superscripts. You can list 3d before or after 4s.
5. Check the superscripts add up to Z.
6. Check for chromium (Z = 24) and copper (Z = 29).

The two exceptions

Atom	Expected from filling order	Actual configuration
Cr, Z = 24	3d⁴ 4s²	1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁵ 4s¹
Cu, Z = 29	3d⁹ 4s²	1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s¹

A half-full or full 3d sub-shell is the usual explanation for the extra stability.

Method in steps: configuration of an ion, 2.2.1(d)(ii)

The specification limits this to s- and p-block ions up to Z = 36. (d-block ions come in 5.3.1.)

1. Write the configuration of the atom.
2. Electrons in the ion = Z − charge (so a 2+ ion has Z − 2, a 3− ion has Z + 3).
3. Cation: remove electrons from the outermost sub-shell first. For Period 4 p-block atoms, take 4p, then 4s. Never take 3d.
4. Anion: add electrons to the outer p-sub-shell until it is full.
5. Check the total.

Small worked reminders

• Up to Z = 36, the heaviest s-block elements are potassium and calcium (rubidium is Z = 37, outside the range).
• Ca²⁺ (Z = 20, 18 electrons): 1s² 2s² 2p⁶ 3s² 3p⁶.
• Ge²⁺ (Z = 32, 30 electrons): remove the two 4p electrons, giving 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s².
• As³⁻ (Z = 33, 36 electrons): 4p fills, giving 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p⁶.

Must-know distinctions

• Shell vs sub-shell vs orbital. A shell contains sub-shells; a sub-shell contains orbitals; an orbital holds a maximum of two electrons.
• Number of orbitals vs number of electrons. A p-sub-shell has 3 orbitals but holds 6 electrons. A d-sub-shell has 5 orbitals but holds 10.
• Filling order vs writing order. 4s fills before 3d, but you may write 3d before 4s when you group by shell. The electrons are the same.
• Atom vs ion. Atoms use Z electrons. Ions use Z minus the charge.
• Paired vs unpaired. An unpaired electron is alone in its orbital. Count them from the box diagram, not from the superscript alone: 3p⁴ has two unpaired electrons, not four.

Where marks are usually lost

• Stating that the fourth shell holds 18 electrons. The answer for 2.2.1(a) is 32.
• Describing an orbital without saying it holds up to two electrons, or without mentioning opposite spins.
• Describing a p-orbital as “figure of eight” without making clear it is a three-dimensional dumbbell with lobes either side of the nucleus.
• In a box diagram, pairing electrons in one p-orbital while another p-orbital in that sub-shell is empty.
• Writing 3d¹ for potassium (Z = 19) because 3d was filled before 4s.
• Writing chromium as 3d⁴ 4s² or copper as 3d⁹ 4s².
• Removing 3d electrons when forming a p-block cation such as Ga³⁺ or Ge²⁺.
• Superscripts that do not add up to the correct electron total, especially for anions.
• Using [Ar] shorthand when the question asks for the full electron configuration.

Quick self-test

1. What is the maximum number of electrons in the second shell?
2. How many orbitals are in a d-sub-shell, and how many electrons can it hold?
3. Describe the shape of a p-orbital.
4. Write the full electron configuration of silicon (Z = 14) and state how many unpaired electrons it has.
5. Write the full electron configuration of cobalt (Z = 27).
6. How many unpaired electrons does a manganese atom (Z = 25) have?
7. Write the electron configuration of K⁺ (Z = 19).
8. Write the electron configuration of N³⁻ (Z = 7).
9. The configuration of an atom is 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p⁵. Identify it.
10. Which ion with a 2− charge has 18 electrons?
11. How many electrons occupy p-orbitals in a chlorine atom (Z = 17)?
12. Why is the configuration of copper 3d¹⁰ 4s¹ rather than 3d⁹ 4s²?

Answers

1. 8 (2s² 2p⁶; 2 × 2² = 8).
2. 5 orbitals, 10 electrons.
3. A dumbbell shape: two lobes on opposite sides of the nucleus. There are three, at right angles to each other.
4. 1s² 2s² 2p⁶ 3s² 3p²; the two 3p electrons are in separate orbitals, so 2 unpaired.
5. 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁷ 4s² (2 + 2 + 6 + 2 + 6 + 7 + 2 = 27).
6. 5 (3d⁵, one electron in each d-orbital).
7. 18 electrons: 1s² 2s² 2p⁶ 3s² 3p⁶.
8. 10 electrons: 1s² 2s² 2p⁶.
9. Total electrons = 35, so it is bromine.
10. Z = 18 − 2 = 16, so S²⁻ (sulfide).
11. 2p⁶ + 3p⁵ = 11.
12. A completely filled 3d sub-shell (d¹⁰) is the more stable, lower-energy arrangement, so one 4s electron occupies 3d instead.

Next steps

• Got a self-test answer wrong? Go back to that section of the study guide.
• Attempt the practice questions, which have fully worked answers.
• Tick off 2.2.1 on the checklist and see the course hub.

Official syllabus

OCR A Level Chemistry A (H432) specification, Version 3.1 (May 2026), for first assessment in 2017, published by OCR. Section 2.2.1 Electron structure.