OCR A Level Chemistry A: The halogens (H432) – Revision Notes

These notes condense section 3.1.3, The halogens, of the OCR A Level Chemistry A (H432) specification, Version 3.1 (May 2026), for first assessment in 2017. They cover outcomes 3.1.3(a) to (g); OCR assesses this module (Module 3, Periodic table and energy) in Periodic table, elements and physical chemistry (01) and Unified chemistry (03). H432 is the full A level. Outcome (g) links to PAG4.

For full explanations and worked examples, use the halogens study guide. When you are ready, move on to the halogens practice questions. The OCR chemistry hub lists every H432 topic, and the printable checklist lets you track each outcome. For a mixed check, try the free diagnostics.

3.1.3(a): Physical properties

• Every halogen exists as diatomic molecules, X₂, joined by one shared pair.
• Solid and liquid halogens are simple molecular lattices.
• Boiling overcomes London forces (induced dipole–dipole interactions) between molecules. The X–X bond stays intact.

	Cl₂	Br₂	I₂
Electrons per molecule	34	70	106
Boiling point (approx.)	−34 °C	59 °C	184 °C
State at room temperature	gas	liquid	solid
Appearance	pale green	red-brown	grey-black (purple vapour)

Method: boiling point trend in three steps

1. Name the force: London forces between diatomic molecules.
2. Compare electron numbers, with the numbers.
3. More electrons, stronger London forces, more energy to overcome them.

3.1.3(b): Electron configuration and redox

• Outer shell s²p⁵; one electron short of a full p sub-shell.
• In many redox reactions: X₂ + 2e⁻ → 2X⁻. Oxidation number 0 → −1, so the halogen is reduced and acts as an oxidising agent.
• Describe redox two ways every time: electron transfer, and oxidation number change. Give the full and the ionic equation.

3.1.3(c): Displacement and colours

Reactivity: Cl₂ > Br₂ > I₂. A halogen displaces only halides below it.

Reaction	Ionic equation	What you see
Cl₂ + Br⁻	Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂	aqueous goes orange; organic layer orange
Cl₂ + I⁻	Cl₂ + 2I⁻ → 2Cl⁻ + I₂	aqueous goes brown; organic layer violet
Br₂ + I⁻	Br₂ + 2I⁻ → 2Br⁻ + I₂	aqueous goes brown; organic layer violet
Br₂ + Cl⁻, I₂ + Cl⁻, I₂ + Br⁻	no reaction	colour of added halogen only

Halogen	In water	In cyclohexane (organic)
Cl₂	pale green	pale green
Br₂	orange	orange
I₂	brown	violet

The organic layer sits on top and shows the colours more clearly. That is why it is added.

3.1.3(d): The reactivity explanation

Down the group:

1. Atomic radius increases (more shells).
2. More shielding by inner electrons.
3. Weaker nuclear attraction for an incoming electron.
4. So a 1− ion forms less easily, and reactivity (oxidising power) decreases.

Point 3 is what links the structure to the trend. Do not leave it out.

3.1.3(e): Disproportionation

Definition: oxidation and reduction of the same element in one reaction.

Reaction	Equation	Chlorine oxidation numbers
Cl₂ with water (water treatment)	Cl₂ + H₂O ⇌ HClO + HCl	0 → +1 (HClO) and 0 → −1 (HCl)
Cl₂ with cold, dilute NaOH (bleach)	Cl₂ + 2NaOH → NaCl + NaClO + H₂O	0 → −1 (NaCl) and 0 → +1 (NaClO)

Names: HClO is chloric(I) acid; NaClO is sodium chlorate(I); ClO⁻ is the chlorate(I) ion.

Analogous reactions, (e)(iii). Swap the halogen or the alkali and keep the pattern.

• Br₂ + H₂O ⇌ HBrO + HBr
• I₂ + 2NaOH → NaI + NaIO + H₂O
• Cl₂ + 2KOH → KCl + KClO + H₂O

Method: proving disproportionation

1. Write the oxidation number of the element in the reactant (0 for an element).
2. Write its oxidation number in each product that contains it.
3. Show one value has gone up (oxidised) and one has gone down (reduced).
4. State that the same element has been oxidised and reduced.

3.1.3(f): Chlorine in water

Benefits	Risks	Other methods
Kills bacteria, preventing diseases such as cholera and typhoid	Chlorine gas is toxic	Ozone
A residual amount keeps water safe in the pipes	Chlorinated hydrocarbons can form from organic matter; some are suspected carcinogens	Ultraviolet light
Cheap and easy to control	Some argue that people lose the right to choose	Filtration (removes solids, not bacteria)

Judgement most often reached: the public health benefit outweighs the small risk. An “evaluate” answer needs that judgement.

3.1.3(g): Halide tests (PAG4)

Method in steps

1. Acidify with dilute nitric acid (removes carbonate, which would also give a precipitate).
2. Add silver nitrate solution; record the colour of the precipitate.
3. Add dilute ammonia, then concentrated ammonia if needed.

Ion	Precipitate	Dilute NH₃	Concentrated NH₃
Cl⁻	white AgCl	soluble	soluble
Br⁻	cream AgBr	insoluble	soluble
I⁻	yellow AgI	insoluble	insoluble

Ionic equation pattern: Ag⁺(aq) + X⁻(aq) → AgX(s)

Ammonia complexes: observations only. In a full anion sequence (3.1.4) test for carbonate, then sulfate, then halide; see qualitative analysis.

Molar masses for halide calculations (Ar: Ag 107.9, Cl 35.5, Br 79.9, I 126.9): AgCl 143.4, AgBr 187.8, AgI 234.8.

Worked reminders

Reminder 1: a displacement answer in full. Chlorine water is added to potassium iodide, then cyclohexane is shaken in.

• See: the aqueous mixture darkens to brown; the top organic layer turns violet.
• Full equation: Cl₂ + 2KI → 2KCl + I₂
• Ionic equation: Cl₂ + 2I⁻ → 2Cl⁻ + I₂
• Redox: chlorine 0 → −1, gains electrons, reduced; iodide −1 → 0, loses electrons, oxidised.

Reminder 2: iodine with water, by analogy. Swap Cl for I in the chlorine–water equation: I₂ + H₂O ⇌ HIO + HI. Iodine goes from 0 to +1 in HIO and from 0 to −1 in HI, so it is disproportionation.

Reminder 3: from precipitate mass to concentration. A white precipitate that dissolves in dilute ammonia forms from 20.0 cm³ of a chloride solution. It weighs 0.287 g.

1. White and soluble in dilute ammonia, so it is AgCl, Mr 143.4.
2. n(AgCl) = 0.287 ÷ 143.4 = 0.00200 mol = n(Cl⁻).
3. c(Cl⁻) = 0.00200 ÷ 0.0200 = 0.100 mol dm⁻³.

Divide the mass by the Mr of the whole silver halide, then by the volume in dm³.

Must-know distinctions

• London forces vs covalent bond: boiling breaks the first, not the second.
• Halogen vs halide: chlorine (Cl₂) is the element; chloride (Cl⁻) is the ion.
• Reduced vs reducing agent: the halogen is reduced; it is the oxidising agent.
• Chlorate(I) vs chlorate(V): ClO⁻ in bleach, ClO₃⁻ only if the question gives a different reaction.
• Dilute vs concentrated ammonia: dilute separates chloride from bromide; concentrated separates bromide from iodide.

Quick self-test

1. Give the full electron configuration of a bromide ion.
2. Why does iodine have a higher boiling point than bromine?
3. What colour is the organic layer when chlorine water is added to potassium bromide and then shaken with cyclohexane?
4. Give the ionic equation when chlorine oxidises iodide ions.
5. Does iodine react with sodium chloride solution? Explain.
6. Define disproportionation.
7. State the oxidation numbers of chlorine in Cl₂, HClO and HCl.
8. Write the equation for iodine with cold, dilute potassium hydroxide.
9. 0.0150 mol of Cl₂ reacts fully with cold, dilute NaOH. What mass of NaClO forms? (Ar: Na 23.0, Cl 35.5, O 16.0)
10. A cream precipitate weighs 0.376 g. How many moles of halide ions were in the sample?
11. Why is hydrochloric acid not used to acidify the sample in a halide test?
12. Give one other method of treating drinking water.

Answers

1. 1s²2s²2p⁶3s²3p⁶3d¹⁰4s²4p⁶
2. I₂ has more electrons (106 against 70), so stronger London forces need more energy to overcome.
3. Orange (bromine has been displaced).
4. Cl₂ + 2I⁻ → 2Cl⁻ + I₂
5. No. Iodine is less reactive than chlorine, so it cannot oxidise Cl⁻.
6. Oxidation and reduction of the same element in the same reaction.
7. 0, +1 and −1.
8. I₂ + 2KOH → KI + KIO + H₂O
9. Mr(NaClO) = 74.5; 0.0150 mol NaClO; 0.0150 × 74.5 = 1.12 g.
10. Cream means AgBr, Mr 187.8; 0.376 ÷ 187.8 = 0.00200 mol Br⁻.
11. HCl brings in its own Cl⁻ ions, so silver chloride would precipitate even with no chloride in the sample.
12. Ozone or ultraviolet light.

Where marks are usually lost

• Saying “bigger molecules” or “more mass” without linking electron number to London force strength.
• Writing that bonds between atoms break on boiling.
• Describing iodine in cyclohexane as brown, or bromine in cyclohexane as violet.
• Writing ionic equations with spectator ions left in, or unbalanced charges such as Cl₂ + I⁻ → Cl⁻ + I₂.
• Giving only “more shielding” for the reactivity trend, with no mention of weaker attraction for the gained electron.
• Proving disproportionation with only one product’s oxidation number.
• Naming NaClO as sodium chlorate without the (I).
• Giving “white” for silver bromide, or claiming AgBr dissolves in dilute ammonia.
• Describing water treatment risks vaguely (“chlorine is dangerous”) instead of naming toxic gas or chlorinated hydrocarbons.
• In mole questions, using the Ar of the halogen instead of the Mr of the silver halide.

Official syllabus

OCR A Level Chemistry A (H432) specification, Version 3.1 (May 2026), for first assessment in 2017, section 3.1.3 The halogens, learning outcomes (a) to (g). Published by OCR.