AQA A-Level Chemistry: Group 7(17), the halogens (7405) – Revision Notes

These notes condense section 3.2.3 Group 7(17), the halogens of the AQA AS and A-level Chemistry specification (7404/7405), version 1.2, July 2026, for AS and A-level exams June 2016 onwards: 3.2.3.1 Trends in properties, 3.2.3.2 Uses of chlorine and chlorate(I), and Required practical 4. All of it is AS content, so nothing here is A-level only. It is examined on AS Paper 1 and A-level Paper 1, and A-level Paper 3 can test any content.

For full explanations and worked examples, read the halogens study guide first. Then test yourself with the halogens practice questions. The course hub is AQA A-Level Chemistry, and the printable checklist lists every specification point. The free 10-minute diagnostics help you find weak spots. For redox basics, see the redox revision notes.

Trends down Group 7 (3.2.3.1)

Property	Trend down the group	Reason
Electronegativity	decreases	bigger atoms, more shielding, bonding pair further from nucleus
Boiling point	increases	more electrons per molecule, stronger van der Waals forces
Oxidising ability of X₂	decreases	incoming electron further from nucleus, more shielding
Reducing ability of X⁻	increases	outer electrons further from nucleus, more shielding, lost more easily
Solubility of AgX in NH₃	decreases	(describe only)

Definition: electronegativity is the power of an atom to attract the pair of electrons in a covalent bond.

Boiling point answer in steps:

1. Simple molecular, diatomic molecules.
2. More electrons per molecule down the group (Cl₂ 34, Br₂ 70, I₂ 106).
3. Stronger van der Waals forces between molecules.
4. More energy needed to overcome them.

Never say the X–X bond breaks on boiling.

Displacement reactions

A halogen displaces any halogen below it from a solution of its halide ions.

Mixture	Result	Ionic equation
Cl₂ + Br⁻	yellow/orange solution	Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂
Cl₂ + I⁻	brown solution	Cl₂ + 2I⁻ → 2Cl⁻ + I₂
Br₂ + I⁻	brown solution	Br₂ + 2I⁻ → 2Br⁻ + I₂
Br₂ + Cl⁻, I₂ + Cl⁻, I₂ + Br⁻	no reaction	–

Halogen = oxidising agent (gains electrons). Halide ion = reducing agent (loses electrons).

Sodium halides with concentrated sulfuric acid

Step 1 for every halide (acid–base, not redox):

NaX + H₂SO₄ → NaHSO₄ + HX       (steamy white fumes of HX)

Halide	Further reaction	Extra observations	S product
F⁻, Cl⁻	none	–	none (+6)
Br⁻	2HBr + H₂SO₄ → Br₂ + SO₂ + 2H₂O	orange-brown fumes	SO₂ (+4)
I⁻	2HI + H₂SO₄ → I₂ + SO₂ + 2H₂O	black solid / purple vapour	SO₂ (+4)
I⁻	6HI + H₂SO₄ → 3I₂ + S + 4H₂O	yellow solid	S (0)
I⁻	8HI + H₂SO₄ → 4I₂ + H₂S + 4H₂O	bad-egg smell	H₂S (−2)

Sulfuric acid half-equations (reductions):

H₂SO₄ + 2H⁺ + 2e⁻ → SO₂ + 2H₂O
H₂SO₄ + 6H⁺ + 6e⁻ → S + 4H₂O
H₂SO₄ + 8H⁺ + 8e⁻ → H₂S + 4H₂O

The lower the oxidation state of sulfur in the product, the stronger the halide’s reducing power.

Silver nitrate test

Method in steps:

1. Add dilute nitric acid to the halide solution.
2. Add silver nitrate solution. Note the precipitate colour.
3. Add dilute ammonia; if the precipitate stays, try concentrated ammonia.

Ion	AgX colour	Dilute NH₃	Conc. NH₃
F⁻	no precipitate	–	–
Cl⁻	white	dissolves	dissolves
Br⁻	cream	no	dissolves
I⁻	yellow	no	no

Three “why” answers the specification names:

• Silver nitrate: forms precipitates of different colours with Cl⁻, Br⁻ and I⁻.
• Acidified (with HNO₃): removes carbonate and hydroxide ions, which would also give precipitates. Not HCl (adds Cl⁻).
• Ammonia: the colours are similar, so solubility in ammonia separates them.

Ag⁺(aq) + Cl⁻(aq) → AgCl(s)
AgCl(s) + 2NH₃(aq) → [Ag(NH₃)₂]⁺(aq) + Cl⁻(aq)

Uses of chlorine and chlorate(I) (3.2.3.2)

Reaction	Equation	Note
Cl₂ + water	Cl₂ + H₂O ⇌ HCl + HClO (or 2H⁺ + Cl⁻ + ClO⁻)	disproportionation, reversible
Cl₂ + water in sunlight	2Cl₂ + 2H₂O → 4H⁺ + 4Cl⁻ + O₂	chloride ions and oxygen
Cl₂ + cold, dilute NaOH	Cl₂ + 2NaOH → NaCl + NaClO + H₂O	disproportionation; NaClO solution is bleach

Oxidation states of chlorine: Cl₂ 0, Cl⁻ −1, ClO⁻ +1.

Water treatment: chlorine kills bacteria. It is toxic and can form chlorinated organic compounds, but the doses are controlled and the health benefits outweigh the toxic effects. Society weighs advantages against disadvantages before adding chemicals (for example sodium fluoride) to water supplies.

Required practical 4 summary

Ion	Reagent	Positive result
Group 2 cations	NaOH(aq); dilute H₂SO₄	precipitate pattern from solubility trends
NH₄⁺	warm with NaOH(aq)	gas turns damp red litmus blue
OH⁻	red litmus	turns blue
CO₃²⁻	dilute acid, then limewater	fizzing; limewater cloudy
SO₄²⁻	acidified BaCl₂(aq)	white precipitate
X⁻	HNO₃, then AgNO₃(aq), then NH₃(aq)	see table above

Small worked reminders

Reminder 1: tracking sulfur with sodium bromide. In 2HBr + H₂SO₄ → Br₂ + SO₂ + 2H₂O, bromine goes from −1 to 0 (oxidised) and sulfur goes from +6 to +4 (reduced). Two bromide ions each lose one electron; one sulfur atom gains two. Electrons lost equal electrons gained, so the equation is consistent.

Reminder 2: building an iodide equation from half-equations. To get iodide reducing H₂SO₄ to SO₂:

H₂SO₄ + 2H⁺ + 2e⁻ → SO₂ + 2H₂O
2I⁻ → I₂ + 2e⁻
Add:  H₂SO₄ + 2H⁺ + 2I⁻ → SO₂ + I₂ + 2H₂O

Electrons cancel because both half-equations already involve two. For S (6e⁻) multiply the iodide half-equation by 3; for H₂S (8e⁻) multiply by 4.

Reminder 3: a silver halide calculation. Excess acidified AgNO₃ gives 1 mol AgX for every 1 mol X⁻. So mass of precipitate ÷ Mr(AgX) = moles of halide ion in the sample. Divide by the volume in dm³ for a concentration. Convert cm³ to dm³ by dividing by 1000.

Reminder 4: why chlorine water is a disproportionation. One Cl atom in each Cl₂ goes up to +1 (in ClO⁻ or HClO) and the other goes down to −1 (in Cl⁻ or HCl). The same element is oxidised and reduced at once.

Reminder 5: order of tests in Required practical 4. If a sample may contain carbonate as well as a halide, test for carbonate first with dilute acid, then use nitric acid before the silver nitrate. Use acidified barium chloride for sulfate so that carbonate does not also give a white precipitate.

Must-know distinctions

• Oxidising ability (halogens) decreases down; reducing ability (halides) increases down.
• Acid–base step (HX made) vs redox step (H₂SO₄ reduced).
• Dilute vs concentrated ammonia: only AgCl dissolves in dilute.
• Chlorine water (equilibrium) vs sunlight (oxygen released).

Quick self-test

1. Define electronegativity.
2. State and explain the trend in electronegativity down Group 7.
3. Which halogen is a liquid at room temperature?
4. Predict the result of adding chlorine water to sodium iodide solution and give the ionic equation.
5. Why does iodine not react with potassium bromide solution?
6. Give the observation when concentrated sulfuric acid is added to solid sodium chloride.
7. What is the oxidation state of sulfur in H₂S, and which halide produces it with concentrated H₂SO₄?
8. A solution gives a cream precipitate with acidified silver nitrate. Which ion is present, and what happens with dilute ammonia?
9. Why is hydrochloric acid not used to acidify silver nitrate?
10. Give the oxidation states of chlorine in Cl₂, Cl⁻ and ClO⁻, and name this type of reaction.
11. Write the equation for chlorine with cold, dilute sodium hydroxide.
12. State one use of the solution formed in question 11.

Answers

1. The power of an atom to attract the pair of electrons in a covalent bond.
2. Decreases; atoms get bigger with more shielding, so the bonding pair is further from the nucleus and less strongly attracted.
3. Bromine.
4. Solution turns brown; Cl₂ + 2I⁻ → 2Cl⁻ + I₂.
5. Iodine is a weaker oxidising agent than bromine, so it cannot oxidise Br⁻.
6. Steamy white fumes (HCl).
7. −2; iodide (sodium iodide).
8. Bromide; the precipitate does not dissolve in dilute ammonia (it dissolves in concentrated).
9. It adds chloride ions, which give a white AgCl precipitate.
10. 0, −1, +1; disproportionation.
11. Cl₂ + 2NaOH → NaCl + NaClO + H₂O.
12. Bleach (also disinfecting or cleaning).

Where marks are usually lost

• Explaining boiling points by covalent bonds breaking.
• Writing “Group 7 gets less reactive because it has more shells” without linking to attraction for an electron.
• Mixing up the two trends: halogens get weaker as oxidising agents, halides stronger as reducing agents.
• Leaving out the acid–base step for NaCl, or calling it redox.
• Giving H₂S as a product with bromide.
• Unbalanced half-equations for sulfuric acid; check charge as well as atoms.
• Writing “AgBr is insoluble in ammonia” without saying dilute.
• Forgetting to name disproportionation for chlorine with water or NaOH.
• Saying chlorine is added to water to “clean” it rather than to kill bacteria.

Official syllabus

AQA, AS and A-level Chemistry specification (7404/7405), version 1.2, July 2026, for AS and A-level exams June 2016 onwards: section 3.2.3 Group 7(17), the halogens (3.2.3.1 Trends in properties and 3.2.3.2 Uses of chlorine and chlorate(I)), including Required practical 4.