Group 17: The Halogens

This guide covers subtopics 11.1 to 11.4 — Physical properties of the Group 17 elements; the chemical properties of the halogen elements and the hydrogen halides; some reactions of the halide ions; and the reactions of chlorine — from Topic 11 of Cambridge International AS & A Level Chemistry 9701, 2025–2027 series. All four are AS Level content, covered together as one continuous story: how the elements behave, how their hydrogen compounds behave, how the ions they form behave, and a closer look at chlorine specifically.

Before studying this

This resource assumes van der Waals forces and volatility trends from Chemical Bonding: Shapes and Intermolecular Forces, and reuses the disproportionation example (chlorine with cold sodium hydroxide) from Redox Processes: Oxidation Numbers and Electron Transfer directly, rather than re-deriving it.

Syllabus coverage

CAMBRIDGE INTERNATIONAL AS & A LEVEL CHEMISTRY 9701 — AS Level, Topic 11

11.1 Physical properties of the Group 17 elements — describing the colours and the trend in volatility of chlorine, bromine and iodine; describing and explaining the trend in bond strength of the halogen molecules; interpreting volatility in terms of instantaneous dipole–induced dipole forces.

11.2 The chemical properties of the halogen elements and the hydrogen halides — describing the relative reactivity of the elements as oxidising agents; describing the reactions of the elements with hydrogen and explaining their relative reactivity; describing the relative thermal stabilities of the hydrogen halides and explaining these in terms of bond strength.

11.3 Some reactions of the halide ions — describing the relative reactivity of halide ions as reducing agents; describing and explaining their reactions with aqueous silver ions followed by aqueous ammonia, and with concentrated sulfuric acid.

11.4 The reactions of chlorine — describing and interpreting, in terms of oxidation number, the reactions of chlorine with cold and with hot aqueous sodium hydroxide as disproportionation reactions; explaining the use of chlorine in water purification via the active species HOCl and ClO⁻.

Physical properties

Halogen	Colour and state at room temperature	Bond energy (kJ mol⁻¹, approx.)
Chlorine, Cl₂	pale yellow-green gas	242
Bromine, Br₂	red-brown liquid	193
Iodine, I₂	grey-black solid (sublimes to a purple vapour)	151

Volatility decreases down the group (Cl₂ is a gas, Br₂ a liquid, I₂ a solid) because volatility is governed by the strength of the instantaneous dipole–induced dipole (van der Waals) forces between molecules — and a larger molecule, with more electrons in a more polarisable electron cloud, forms stronger van der Waals forces, needing more energy to separate.

Bond strength within each X–X molecule also decreases down the group — 242 → 193 → 151 kJ mol⁻¹ — because the bonding atoms get larger, so the shared electron pair sits further from both nuclei, in a longer, weaker bond. It’s worth noticing that these two trends run in the same direction but for unrelated reasons: one is about forces between separate molecules, the other about the bond within a single molecule.

Reactivity of the elements as oxidising agents

Oxidising power decreases down the group: Cl₂ is the strongest oxidising agent of the three, I₂ the weakest. A larger halogen atom holds an incoming electron less tightly (further from the nucleus, more shielding), so it’s less able to oxidise another species by taking an electron from it.

This shows up directly in reactions with hydrogen: H₂ + Cl₂ → 2HCl proceeds explosively in sunlight; H₂ + Br₂ → 2HBr needs heating and a catalyst and is markedly less vigorous; H₂ + I₂ ⇌ 2HI is slow, needs heat and a catalyst, and — notably — doesn’t even go to completion, remaining an equilibrium mixture.

Thermal stability of the hydrogen halides follows the same trend: HCl is stable to heat, HBr less so, and HI decomposes significantly on strong heating. The explanation is bond strength: the H–X bond weakens down the group (a larger halogen atom gives a longer, weaker bond to the small hydrogen atom), so less energy is needed to break it apart again.

Reactivity of the halide ions as reducing agents

This is the reverse trend to the elements’ oxidising power, and it’s important not to confuse the two: reducing power of the halide ions increases down the group — I⁻ is the strongest reducing agent, Cl⁻ the weakest. A larger halide ion holds its outer electrons less tightly, so it gives one up (is oxidised) more readily.

Silver nitrate, then aqueous ammonia — a classic identification sequence. Each halide ion gives a different-coloured silver halide precipitate, and the precipitates differ in how readily they redissolve in ammonia (reflecting decreasing solubility, i.e. an increasingly strong Ag–X interaction, down the group):

Halide	Precipitate with Ag⁺(aq)	Dissolves in dilute NH₃(aq)?	Dissolves in concentrated NH₃(aq)?
Cl⁻	white	yes	yes
Br⁻	cream	no	yes
I⁻	yellow	no	no

Concentrated sulfuric acid — the most information-rich test, because it shows the reducing-power trend directly through how far each halide reduces sulfur. Every halide first undergoes a simple acid–base (ligand-exchange) step giving misty/steamy fumes of the hydrogen halide:

NaX + H₂SO₄ → NaHSO₄ + HX

Chloride stops there — Cl⁻ isn’t a strong enough reducing agent to reduce sulfur, so only steamy HCl fumes are seen. Bromide goes further, partially reducing sulfur from +6 to +4:

2HBr + H₂SO₄ → Br₂ + SO₂ + 2H₂O

giving steamy fumes plus orange/brown bromine vapour and choking SO₂ gas. Iodide, the strongest reducing agent, reduces sulfur furthest, down to −2:

8HI + H₂SO₄ → 4I₂ + H₂S + 4H₂O

giving steamy fumes, a black/purple solid (or violet vapour) of iodine, and the rotten-egg smell of H₂S.

Worked example. A student is given three unlabelled solid samples, each containing a single sodium halide, and adds concentrated sulfuric acid to each. Sample A gives only steamy fumes. Sample B gives steamy fumes and a smell of rotten eggs. Identify A and B.

Steamy fumes alone, with no further reaction, indicates the halide is too weak a reducing agent to reduce sulfur at all — this is the behaviour of chloride. Sample A is NaCl. The rotten-egg smell is H₂S, meaning sulfur has been reduced all the way from +6 to −2 — only the strongest reducing agent, iodide, goes this far. Sample B is NaI.

The reactions of chlorine

Cold, dilute sodium hydroxide — a disproportionation already covered in Redox Processes:

Cl₂ + 2NaOH → NaCl + NaOCl + H₂O

Hot, concentrated sodium hydroxide gives a different disproportionation, with chlorine oxidised further, to chlorate(V) instead of chlorate(I):

3Cl₂ + 6NaOH → 5NaCl + NaClO₃ + 3H₂O

In both cases chlorine (oxidation number 0) is simultaneously reduced to −1 (in NaCl) and oxidised — to +1 in the cold reaction, +5 in the hot one.

Water purification: chlorine reacts with water to form a mixture containing the active disinfecting species, HOCl (chloric(I) acid) and the chlorate(I) ion, ClO⁻:

Cl₂ + H₂O ⇌ HOCl + HCl

Both HOCl and ClO⁻ are strong enough oxidising agents to kill disease-causing bacteria in the water supply, which is the practical reason chlorine is added to drinking water.

Common mistakes

• Mixing up “oxidising power of the elements” (decreases down the group) with “reducing power of the ions” (increases down the group). They’re opposite trends describing different species — write down which one a question is actually asking about before answering.
• Forgetting that bromide only partially reduces sulfur, while iodide goes all the way to H₂S. The concentrated sulfuric acid test is only useful as an identification test because the three halides stop at different points.
• Writing the hot/concentrated chlorine–NaOH reaction with the same products as the cold/dilute one. Cold dilute gives chlorate(I), NaOCl; hot concentrated gives chlorate(V), NaClO₃ — different oxidation numbers, different balanced equations.
• Describing AgCl as “insoluble in ammonia” by confusing it with AgBr or AgI. Only AgCl dissolves in dilute ammonia — make sure to state which concentration of ammonia is being used.

Quick revision checklist

• Colours, states and volatility trend of Cl₂, Br₂, I₂; bond strength decreasing down the group
• Oxidising power of the elements decreases down the group; reflected in reactions with hydrogen and thermal stability of the hydrogen halides
• Reducing power of the halide ions increases down the group
• AgNO₃/NH₃ test: precipitate colours and ammonia solubility
• Concentrated H₂SO₄ test: Cl⁻ (fumes only), Br⁻ (fumes + SO₂ + Br₂), I⁻ (fumes + H₂S + I₂)
• Chlorine + cold dilute NaOH → chlorate(I); chlorine + hot concentrated NaOH → chlorate(V); both disproportionations
• HOCl and ClO⁻ as the active species in water purification

Related resources

• Redox Processes: Oxidation Numbers and Electron Transfer — the disproportionation framework this resource builds on
• Group 2: The Alkaline Earth Metals — a second group studied the same way
• Identification of Ions and Gases — the IGCSE/O Level halide and gas tests this resource extends
• Cambridge AS & A Level Chemistry hub

Written against Cambridge International AS & A Level Chemistry 9701, 2025–2027 series. Always check the current syllabus for your examination year.