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

This guide teaches 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. It covers both sub-sections: 3.2.3.1 Trends in properties and 3.2.3.2 Uses of chlorine and chlorate(I), plus Required practical 4, which the specification lists in this section. All of it is AS content, so nothing here is A-level only. The specification places 3.2.3 on AS Paper 1 (inorganic sections 3.2.1 to 3.2.3) and on A-level Paper 1 (all of section 3.2), and A-level Paper 3 can test any content.

For quick recall, use the halogens revision notes. To test yourself, try the halogens practice questions. The course hub is AQA A-Level Chemistry, the printable checklist lists every specification point, and the free 10-minute diagnostics show where your gaps are. Oxidation states and half-equations are taught in oxidation, reduction and redox equations, and van der Waals forces in bonding.

What this topic covers

Spec section	What you must be able to do
3.2.3.1	Explain the trend in electronegativity down the group
3.2.3.1	Explain the trend in boiling point of the elements from structure and bonding
3.2.3.1	Describe the trend in oxidising ability, including displacement of halide ions in aqueous solution
3.2.3.1	Describe the trend in reducing ability of halide ions, including solid sodium halides with concentrated sulfuric acid
3.2.3.1	Use acidified silver nitrate to identify halide ions; explain why it is used, why it is acidified and why ammonia is added
3.2.3.1	Describe the trend in solubility of the silver halides in ammonia
3.2.3.2	Chlorine with water: to chloride and chlorate(I) ions, and to chloride ions and oxygen
3.2.3.2	Chlorine in water treatment; weighing benefits against toxic effects
3.2.3.2	Chlorine with cold, dilute, aqueous NaOH, and uses of the solution
Required practical 4	Test-tube tests for Group 2 cations, NH₄⁺, halide ions, OH⁻, CO₃²⁻ and SO₄²⁻

The specification notes that fluorine is too dangerous for a school laboratory, so practical work uses chlorine, bromine and iodine.

Electronegativity

Electronegativity is the power of an atom to attract the pair of electrons in a covalent bond. It decreases down Group 7, so fluorine is the most electronegative element in the group.

Going down the group:

1. Atomic radius increases, because each element has one more shell of electrons.
2. There is more shielding from inner shells.
3. So the bonding pair is further from the nucleus and is attracted less strongly, even though the nuclear charge increases.

Boiling points

Element	Formula	State at room temperature	Electrons per molecule	Boiling point / °C
Fluorine	F₂	pale yellow gas	18	−188
Chlorine	Cl₂	pale green gas	34	−34
Bromine	Br₂	red-brown liquid	70	59
Iodine	I₂	grey-black solid	106	184

Boiling point increases down the group.

Worked example 1: explaining the boiling point trend

Explain why iodine has a higher boiling point than chlorine.

• Both are simple molecular substances made of diatomic molecules.
• I₂ has more electrons (106) than Cl₂ (34), so I₂ molecules have larger temporary dipoles.
• The van der Waals forces between I₂ molecules are therefore stronger.
• More energy is needed to overcome them, so iodine boils at a higher temperature.

The I–I covalent bond is not broken when iodine boils. Only the forces between molecules are overcome.

Oxidising ability and displacement

A halogen acts as an oxidising agent by gaining electrons: X₂ + 2e⁻ → 2X⁻. Oxidising ability decreases down the group. The atoms get bigger and shielding increases, so an incoming electron is attracted less strongly.

A more reactive halogen displaces a less reactive one from a solution of its halide ions.

Halogen added	KCl(aq)	KBr(aq)	KI(aq)
Cl₂(aq)	no change	turns yellow/orange (Br₂)	turns brown (I₂)
Br₂(aq)	no change	no change	turns brown (I₂)
I₂(aq)	no change	no change	no change

Worked example 2: bromine water and potassium iodide

Bromine water is added to potassium iodide solution. Give the observation, the ionic equation and the half-equations.

Observation:      colourless solution turns brown
Ionic equation:   Br₂ + 2I⁻ → 2Br⁻ + I₂
Reduction:        Br₂ + 2e⁻ → 2Br⁻
Oxidation:        2I⁻ → I₂ + 2e⁻

Bromine is the oxidising agent, iodide ions are the reducing agent and potassium ions are spectators. Iodine cannot oxidise bromide ions, so I₂ added to KBr(aq) shows no reaction.

Reducing ability of halide ions

A halide ion acts as a reducing agent by losing electrons: 2X⁻ → X₂ + 2e⁻. Reducing ability increases down the group: F⁻ < Cl⁻ < Br⁻ < I⁻. A larger ion holds its outer electrons further from the nucleus with more shielding, so it loses an electron more easily.

Concentrated sulfuric acid shows the trend. Every solid sodium halide first reacts in an acid–base reaction to give the hydrogen halide:

NaCl + H₂SO₄ → NaHSO₄ + HCl

Halide	What happens	Observations	Sulfur product (oxidation state)
NaF, NaCl	acid–base only; HF or HCl formed	steamy white fumes	none, S stays +6
NaBr	HBr formed, then 2HBr + H₂SO₄ → Br₂ + SO₂ + 2H₂O	steamy fumes, orange-brown fumes of Br₂	SO₂ (+4)
NaI	HI formed, then H₂SO₄ is reduced to SO₂, S and H₂S	steamy fumes, black solid or purple vapour of I₂, yellow solid (S), smell of bad eggs (H₂S)	SO₂ (+4), S (0), H₂S (−2)

The further the oxidation state of sulfur falls, the stronger the reducing agent. Iodide ions reduce sulfur all the way from +6 to −2.

Worked example 3: iodide reducing sulfuric acid to hydrogen sulfide

Write a half-equation for H₂SO₄ forming H₂S, and an overall ionic equation for its reaction with iodide ions.

Step 1  balance S:               H₂SO₄ → H₂S
Step 2  balance O with H₂O:      H₂SO₄ → H₂S + 4H₂O
Step 3  balance H with H⁺:       H₂SO₄ + 8H⁺ → H₂S + 4H₂O
Step 4  balance charge with e⁻:  H₂SO₄ + 8H⁺ + 8e⁻ → H₂S + 4H₂O
Step 5  iodide half-equation ×4: 8I⁻ → 4I₂ + 8e⁻
Step 6  add:                     H₂SO₄ + 8H⁺ + 8I⁻ → H₂S + 4I₂ + 4H₂O

Check: sulfur goes from +6 to −2, a gain of 8 electrons, matching 8 iodide ions each losing one electron.

Identifying halide ions with silver nitrate

Silver ions form precipitates with chloride, bromide and iodide ions, and the colours differ, which is why silver nitrate solution is used:

Ag⁺(aq) + X⁻(aq) → AgX(s)

Ion	Precipitate	Dilute NH₃(aq)	Concentrated NH₃(aq)
F⁻	none (AgF is soluble)	not needed	not needed
Cl⁻	white AgCl	dissolves	dissolves
Br⁻	cream AgBr	does not dissolve	dissolves
I⁻	yellow AgI	does not dissolve	does not dissolve

Solubility of the silver halides in ammonia decreases from AgCl to AgI.

Why acidify? Add dilute nitric acid first to react with any carbonate or hydroxide ions. Otherwise they would form other silver precipitates (such as silver carbonate) and give a false result. Do not use hydrochloric acid, because it adds chloride ions.

Why add ammonia? White, cream and yellow are hard to tell apart, especially in small amounts. Ammonia confirms the halide by solubility. The dissolving forms a complex ion, for example:

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

Worked example 4: identifying a halide by mass

0.300 g of a potassium halide, KX, is dissolved and treated with excess acidified silver nitrate. The dried precipitate of AgX has a mass of 0.473 g. Identify X. (Ar: K = 39.1, Ag = 107.9)

Moles KX = moles AgX
0.300 / (39.1 + x) = 0.473 / (107.9 + x)
0.300(107.9 + x) = 0.473(39.1 + x)
32.37 + 0.300x = 18.494 + 0.473x
13.876 = 0.173x
x = 80.2

Ar(X) ≈ 80, so X is bromine, and the precipitate would be cream.

Chlorine and water

Chlorine reacts with water in a reversible reaction to form chloride and chlorate(I) ions:

Cl₂ + H₂O ⇌ HCl + HClO      or      Cl₂ + H₂O ⇌ 2H⁺ + Cl⁻ + ClO⁻

This is disproportionation: chlorine is oxidised (0 to +1 in ClO⁻) and reduced (0 to −1 in Cl⁻) in the same reaction.

In bright sunlight, chlorine water decomposes to chloride ions and oxygen:

2Cl₂ + 2H₂O → 4H⁺ + 4Cl⁻ + O₂

Chlorine in water treatment

Chlorine is added to drinking water to kill bacteria. Chlorine is toxic, and it can react with organic matter in water to form chlorinated organic compounds that may be harmful. The amounts used are small and controlled. The specification asks you to appreciate that the health benefits of chlorination outweigh its toxic effects, because it prevents water-borne diseases.

More generally, society weighs advantages and disadvantages before adding any chemical to a water supply. Sodium fluoride is one example: it is added in some areas to reduce tooth decay, but people disagree over adding a substance to everyone’s water without individual consent.

Chlorine and cold, dilute sodium hydroxide

Cl₂ + 2NaOH → NaCl + NaClO + H₂O
Cl₂ + 2OH⁻ → Cl⁻ + ClO⁻ + H₂O

This is also disproportionation. The solution contains sodium chlorate(I), NaClO, and is used as bleach, for disinfecting and cleaning.

Worked example 5: mass of sodium chlorate(I)

14.2 g of chlorine reacts completely with cold, dilute sodium hydroxide. Calculate the maximum mass of NaClO formed. (Mr: Cl₂ = 71.0, NaClO = 74.5)

n(Cl₂)  = 14.2 / 71.0 = 0.200 mol
n(NaClO) = n(Cl₂) = 0.200 mol      (1 : 1 ratio)
mass    = 0.200 × 74.5 = 14.9 g

Required practical 4

Ion	Test	Positive result
Group 2 cations	add NaOH(aq), and separately dilute H₂SO₄	compare precipitates using the Group 2 hydroxide and sulfate solubility trends
NH₄⁺	warm with NaOH(aq); hold damp red litmus at the mouth of the tube	litmus turns blue (ammonia)
OH⁻	red litmus or universal indicator	blue litmus; indicator shows alkaline
CO₃²⁻	add dilute acid; bubble gas through limewater	effervescence; limewater turns cloudy
SO₄²⁻	add acidified BaCl₂(aq)	white precipitate
Halide ions	add dilute HNO₃ then AgNO₃(aq); then NH₃(aq)	white, cream or yellow precipitate, as above

Common errors

• Saying covalent bonds break when halogens boil.
• Comparing van der Waals forces by “size” without linking to the number of electrons.
• Writing that bromide reduces sulfuric acid to H₂S; only iodide reaches H₂S.
• Calling the NaCl reaction redox. No oxidation state changes.
• Acidifying with HCl before the silver nitrate test.
• Saying AgBr is insoluble in all ammonia. It dissolves in concentrated ammonia.
• Missing the word “disproportionation” when chlorine forms Cl⁻ and ClO⁻.

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.