OCR A Level Chemistry A: Electron structure (H432)

This study guide teaches 2.2.1 Electron structure, part of Module 2 (Foundations in chemistry) of the OCR A Level Chemistry A (H432) specification, Version 3.1 (May 2026), for first assessment in 2017. All of 2.2.1(a) to (d) is taught here, along with the specification’s additional guidance. There are no tiers in H432, so every candidate needs all of this content. Module 2 content can be examined in any of the three written components, 01, 02 and 03.

When you have worked through it, use the electron structure revision notes for quick recall and the electron structure practice questions to test yourself. For the wider course, go to the H432 course hub; to record progress statement by statement, print the course checklist. A free diagnostic is a quick way to find which ideas need most work.

What this unit covers

Specification reference	What you must be able to do
2.2.1(a)	State how many electrons can fill each of the first four shells
2.2.1(b)(i)	Describe an atomic orbital as a region around the nucleus holding up to two electrons with opposite spins
2.2.1(b)(ii)	Describe the shapes of s- and p-orbitals
2.2.1(b)(iii)	Give the number of orbitals in s-, p- and d-sub-shells and how many electrons each sub-shell holds
2.2.1(c)(i)	Fill orbitals for the first three shells and the 4s and 4p orbitals in order of increasing energy
2.2.1(c)(ii)	Place electrons singly in orbitals of the same energy before pairing them
2.2.1(d)(i)	Deduce the electron configuration of any atom up to Z = 36 from its atomic number
2.2.1(d)(ii)	Deduce the configuration of s- and p-block ions up to Z = 36 from atomic number and charge
Additional guidance	Use sub-shell notation (the specification’s example is oxygen, 1s²2s²2p⁴) and “electrons in box” diagrams; HSW1 and HSW7 on how models of electron structure developed

No practical activity group (PAG) is linked to 2.2.1 in the specification. Before you start, make sure you are secure on protons, neutrons, electrons and atomic number from atomic structure and isotopes.

Why the model has layers (HSW1, HSW7)

At GCSE you probably drew electrons in shells of 2, 8, 8. That model is useful, but it breaks down after argon. It cannot say why potassium’s 19th electron goes into the fourth shell while the third shell still has room. It also cannot explain the small drops in first ionisation energy partway across a period, which you will meet in 3.1.1.

The fix was to refine the model. Each shell is split into sub-shells (s, p, d), and each sub-shell is split into orbitals. Sub-shells have different energies, and electrons fill the lowest energy available. This is a good example of a scientific model being improved when new evidence does not fit the old one, which is the point the specification’s HSW references make.

Shells and how many electrons they hold: 2.2.1(a)

The shell number is often written n. The maximum number of electrons in shell n is 2n².

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

You only need s, p and d sub-shells in detail for 2.2.1. The fourth shell’s total of 32 includes a 4f sub-shell (seven orbitals, 14 electrons), which is why the number is larger than s + p + d alone would give.

Worked example 1: counting the third shell

Show that the third shell can hold 18 electrons.

1. The third shell contains 3s, 3p and 3d sub-shells.
2. Orbitals: 3s has 1, 3p has 3, 3d has 5. Total = 9 orbitals.
3. Each orbital holds a maximum of 2 electrons, so 9 × 2 = 18 electrons.
4. Check with 2n²: 2 × 3² = 18.

Atomic orbitals: 2.2.1(b)

What an orbital is

An atomic orbital is a region around the nucleus that can hold up to two electrons, with opposite spins. “Opposite spins” is why you draw a paired orbital as one up arrow and one down arrow. Two electrons in the same orbital can never have the same spin.

An orbital is not a fixed path like a planet’s orbit. It describes where an electron is likely to be found.

Shapes of s- and p-orbitals

• s-orbital: spherical, centred on the nucleus. Each shell has one s-orbital, and the s-orbital of a higher shell is a larger sphere (2s is bigger than 1s).
• p-orbital: dumbbell-shaped, with two lobes on opposite sides of the nucleus. A p-sub-shell has three p-orbitals at right angles to each other, along the x, y and z axes. They are often labelled px, py and pz.

When you sketch a p-orbital, draw the two lobes meeting at the nucleus, and draw three of them on perpendicular axes if asked for the whole sub-shell. You are not asked for d-orbital shapes in 2.2.1.

Orbitals and electrons in each sub-shell

Sub-shell type	Orbitals it contains	Electrons it can hold
s	1	2
p	3	6
d	5	10

Remember the pattern: orbitals go 1, 3, 5, and electrons are always double the number of orbitals.

Filling orbitals: 2.2.1(c)

Order of increasing energy

Electrons go into the lowest-energy orbital available. For the first three shells plus 4s and 4p, the order is:

1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p

The key point is that 4s is filled before 3d. This is why potassium and calcium put their outer electrons into 4s, and why the d-block only begins at scandium (Z = 21).

Singly before pairing

Orbitals in the same sub-shell have the same energy. When electrons enter a p- or d-sub-shell, they occupy the orbitals singly first, with parallel spins, and only pair up once every orbital in that sub-shell has one electron. Electrons repel each other, so they spread out between orbitals of equal energy before sharing one.

Electrons in boxes

In an “electrons in box” diagram, each box is one orbital and each arrow is one electron. A sub-shell is drawn as a group of joined boxes: one for s, three for p, five for d.

Electron configurations of atoms: 2.2.1(d)(i)

Use sub-shell notation: the shell number, the sub-shell letter and the number of electrons as a superscript. Fill in the energy order above until you have used up all Z electrons.

Worked example 2: arsenic, Z = 33

Write the full electron configuration of arsenic and draw its outer sub-shells as boxes.

1. Fill in energy order: 1s² (2), 2s² (4), 2p⁶ (10), 3s² (12), 3p⁶ (18), 4s² (20), 3d¹⁰ (30), 4p³ (33).
2. Written in shell order: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p³.
3. Check: 2 + 2 + 6 + 2 + 6 + 10 + 2 + 3 = 33.
4. Boxes for the fourth-shell and 3d electrons:

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

The three 4p electrons go into separate orbitals with parallel spins, so arsenic has three unpaired electrons.

You can write the 3d sub-shell before or after 4s. Writing 3d¹⁰ 4s² groups the third-shell electrons together; writing 4s² 3d¹⁰ follows filling order. Both describe the same atom. The specification’s own example for iron in 5.3.1 lists 3d before 4s.

Worked example 3: vanadium, Z = 23

1. After argon’s 18 electrons (1s² 2s² 2p⁶ 3s² 3p⁶), five remain.
2. Two go into 4s, then three into 3d.
3. Configuration: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d³ 4s².
4. The three 3d electrons occupy three of the five d-orbitals singly, so vanadium has three unpaired electrons.

Chromium and copper

Two atoms up to Z = 36 do not follow the simple filling order:

• Chromium (Z = 24): 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁵ 4s¹, not 3d⁴ 4s².
• Copper (Z = 29): 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s¹, not 3d⁹ 4s².

In each case one 4s electron moves into 3d. The usual explanation is that a half-filled (d⁵) or completely filled (d¹⁰) d-sub-shell gives a lower-energy, more stable arrangement. Learn these two as facts; every other atom up to krypton follows the order.

Shorthand

You may see [Ar] used for 1s² 2s² 2p⁶ 3s² 3p⁶, so vanadium becomes [Ar] 3d³ 4s². The specification asks for sub-shell notation, so when a question asks for the full configuration, write every sub-shell out.

Electron configurations of ions: 2.2.1(d)(ii)

For 2.2.1, ions are limited to the s- and p-blocks up to Z = 36. Ions of d-block elements are taught in 5.3.1.

• Positive ion: start from the atom, then remove electrons from the highest-energy outer sub-shell first. For a p-block element in Period 4, that means 4p first, then 4s.
• Negative ion: start from the atom and add electrons to the next available orbitals, normally completing the outer p-sub-shell.

Worked example 4: gallium ion, Ga³⁺ (Z = 31)

1. Gallium atom: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p¹ (31 electrons).
2. Ga³⁺ has 31 − 3 = 28 electrons.
3. Remove the 4p electron, then both 4s electrons.
4. Ga³⁺: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰.

The filled 3d sub-shell stays. Taking 3d electrons and leaving 4s² is the mistake to avoid.

Worked example 5: selenide ion, Se²⁻ (Z = 34)

1. Selenium atom: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p⁴.
2. Se²⁻ has 34 + 2 = 36 electrons.
3. The two extra electrons complete 4p.
4. Se²⁻: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p⁶, the same as a krypton atom.

Species with the same electron configuration are called isoelectronic. Se²⁻ and Kr are an example; so are Al³⁺ and Ne, which both have 1s² 2s² 2p⁶.

Common errors

• Giving the fourth shell a maximum of 18 because only s, p and d were counted. It is 32.
• Defining an orbital as “a shell” or “an orbit”. Use the specification’s idea: a region that holds up to two electrons with opposite spins.
• Drawing two electrons in one p-orbital while another p-orbital in the same sub-shell is still empty.
• Drawing paired electrons with both arrows pointing the same way.
• Writing 3d before 4s in the filling order when deducing a configuration, which puts potassium’s outer electron in 3d.
• Forgetting the chromium and copper exceptions, or applying a similar “exception” to other atoms.
• For a Period 4 p-block cation, removing 3d electrons. Remove 4p, then 4s.
• Not checking that the superscripts add up to the number of electrons (Z for an atom, adjusted for charge in an ion).

Where this leads

2.2.1 feeds directly into 2.2.2 (bonding and structure), where outer electrons explain ionic and covalent bonding, and into 3.1.1, where configurations explain the s-, p- and d-blocks and ionisation energy trends. In 5.3.1 you return to configurations for d-block atoms and ions.

Next steps

• Recall: electron structure revision notes
• Practise: electron structure practice questions
• Course map: H432 course hub and checklist
• Exam technique: preparing for the OCR chemistry exams
• Find your gaps: free diagnostics

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.