HookThe gas that killed at Ypres keeps your tap water safe
On 22 April 1915, near Ypres, German engineers opened around 5,700 pressurised cylinders and let roughly 168 tonnes of chlorine drift as a green-yellow cloud into the Allied trenches — the first mass chemical attack in history. Yet in the same decade, the same gas was quietly running one of the greatest life-saving interventions ever attempted. Lincoln's waterworks began dosing its supply with chlorine in 1905 after a typhoid epidemic; Jersey City followed with continuous chlorination in 1908; and cities that had watched typhoid kill around 25 in every 100,000 residents a year in 1900 saw the disease all but vanish by mid-century. One element, two ledgers — and AQA writes the tension straight into the specification: you are expected to weigh chlorine's toxicity against the benefit of drinking water that does not carry cholera and typhoid.
Group 7 is the mirror image of the Group 2 story you have just finished. Instead of metals giving electrons away ever more easily down the group, these are non-metals whose grip on incoming electrons weakens with every added shell. Four stories carry the section: the trends in electronegativity, boiling point and oxidising power; the halide ions as reducing agents, revealed by the concentrated sulfuric acid ladder; the silver nitrate and ammonia tests that identify which halide you hold; and the chemistry of chlorine in water, where one element manages to oxidise and reduce itself in the same reaction. Required Practical 4 then packages the tests into a single method.
DataThree trends, one cause — the outer shell moves further away
Electronegativity falls down the group: fluorine 4.0, chlorine 3.0, bromine 2.8, iodine 2.5 on the Pauling scale. A bonding pair of electrons sits further from an iodine nucleus than a fluorine one, behind more shells of shielding, so the pull on it is weaker — fluorine is the most electronegative element there is.
Boiling point rises down the group, and the colours rise with it: fluorine is a pale yellow gas (−188 °C), chlorine a green gas (−34 °C), bromine a red-brown liquid (59 °C), iodine a grey-black solid whose vapour is violet (184 °C). The molecules are all non-polar X₂, so the only forces between them are van der Waals induced-dipole attractions — and those scale with the number of electrons. I₂ carries 106 electrons to F₂'s 18, so its molecules stick together into a solid at room temperature. Keep this trend in its own drawer: students fuse 'reactivity falls' with 'boiling point falls' and get the physical trend backwards.
Oxidising power falls down the group. A halogen oxidises by capturing an electron into its outer shell; down the group that shell is further out and more shielded, so the capture is less favourable. The evidence is displacement: chlorine water added to potassium bromide solution turns it orange — \(\text{Cl}_2 + 2\text{Br}^- \rightarrow 2\text{Cl}^- + \text{Br}_2\) — and added to potassium iodide turns it brown as iodine forms. Bromine displaces iodide but cannot touch chloride. Each halogen displaces every halide below it and none above it: a strict pecking order, readable straight off a results table.
MechanismThe sulfuric acid ladder — how far can each halide push?
Flip the oxidising trend over and you get the section's most examined idea: halide ions get better at reducing down the group. A halide reduces by giving its extra electron away, and the bigger the ion — outer electron further from the nucleus, more shielded — the more easily that electron leaves. Iodide, the giant, is the strongest reducing agent of the set; chloride is feeble.
The test bench is drops of concentrated sulfuric acid on the solid sodium halide, and each halide climbs as far up the reduction ladder as its strength allows. Sodium chloride: misty white fumes of HCl and nothing more — \(\text{NaCl} + \text{H}_2\text{SO}_4 \rightarrow \text{NaHSO}_4 + \text{HCl}\). Check the oxidation states: nothing changes. This is an acid–base proton transfer, because chloride cannot reduce sulfur from +6 at all. Sodium bromide: misty HBr fumes first, then brown Br₂ vapour and colourless, choking SO₂ — bromide manages a two-electron push, \(2\text{HBr} + \text{H}_2\text{SO}_4 \rightarrow \text{Br}_2 + \text{SO}_2 + 2\text{H}_2\text{O}\), taking sulfur from +6 to +4. Sodium iodide: the full firework show — violet fumes and grey-black solid iodine, SO₂, pale yellow specks of solid sulfur, and the rotten-egg smell of H₂S. Iodide drags sulfur down the entire staircase, +6 → +4 → 0 → −2.
Every product is an observation, and AQA asks for them as observations: 'misty fumes', 'brown gas', 'purple vapour', 'yellow solid', 'gas smelling of bad eggs' — each one a mark, each one evidence of how much reducing power the halide brought to the fight.
The electron audit that checks any equation on this ladder — take iodide's deepest push: \(8\text{HI} + \text{H}_2\text{SO}_4 \rightarrow 4\text{I}_2 + \text{H}_2\text{S} + 4\text{H}_2\text{O}\). Oxidation half-equation: \(2\text{I}^- \rightarrow \text{I}_2 + 2\text{e}^-\), running four times = 8 electrons released. Reduction half-equation: \(\text{H}_2\text{SO}_4 + 8\text{H}^+ + 8\text{e}^- \rightarrow \text{H}_2\text{S} + 4\text{H}_2\text{O}\) — sulfur falls from +6 to −2, a gain of exactly 8 electrons. Electrons released = electrons gained, and the atoms balance: 8 iodine, 1 sulfur, 4 oxygen, 10 hydrogen each side. The same audit on the bromide equation shows 2 electrons each way. If your electron books don't balance, the equation is wrong — the audit takes twenty seconds and catches almost every slip before the examiner does.
DataSilver nitrate, then ammonia — a two-stage sieve for halides
To identify a halide in solution, first acidify with dilute nitric acid. The acid destroys carbonate and hydroxide ions, which would otherwise precipitate with silver and hand you a white lie. The choice of acid is itself examinable: hydrochloric acid would add the very chloride ions you are testing for, and sulfuric acid would precipitate silver sulfate — nitric is the only safe option.
Then add silver nitrate solution: \(\text{Ag}^+\text{(aq)} + \text{X}^-\text{(aq)} \rightarrow \text{AgX(s)}\). Chloride gives a white precipitate, bromide cream, iodide yellow. Fluoride gives no precipitate at all — silver fluoride is soluble — which is itself the identification. The colours, though, are a judgement call in a test tube: cream against white against pale yellow under fluorescent light is not a distinction you want your grade riding on.
So the test has a second stage. Dilute ammonia redissolves silver chloride only. Concentrated ammonia redissolves silver bromide as well. Silver iodide defeats both, sitting stubborn and yellow at the bottom of the tube. The solubility of the silver halides falls down the group, and the ammonia ladder converts a colour judgement into a clean yes/no pair: dissolves in dilute = chloride; only in concentrated = bromide; in neither = iodide.
MechanismChlorine in water — one element, both directions at once
Bubble chlorine into water and it reacts with itself: \(\text{Cl}_2 + \text{H}_2\text{O} \rightleftharpoons \text{HClO} + \text{HCl}\). Track the oxidation states and something strange has happened: chlorine enters at 0 and leaves as both +1 (in chloric(I) acid, HClO) and −1 (in HCl). One element, simultaneously oxidised and reduced, in the same reaction — a disproportionation. The +1 species is the worker: chloric(I) acid and the chlorate(I) ion ClO⁻ are oxidising agents aggressive enough to kill bacteria, which is the entire point of putting chlorine in drinking water and swimming pools.
Sunlight runs a different, wasteful reaction: \(2\text{Cl}_2 + 2\text{H}_2\text{O} \rightarrow 4\text{H}^+ + 4\text{Cl}^- + \text{O}_2\) — chlorine is lost as chloride while the water is oxidised to oxygen, which is why outdoor pools need constant re-dosing on bright days. And run the disproportionation in cold, dilute sodium hydroxide instead of water and you manufacture household bleach: \(\text{Cl}_2 + 2\text{NaOH} \rightarrow \text{NaCl} + \text{NaClO} + \text{H}_2\text{O}\), the mixture of chloride (−1) and sodium chlorate(I) (+1).
The judgement AQA asks for is the Ypres-versus-typhoid balance from the top of the page: chlorine is toxic, and traces of chlorinated organic by-products in treated water carry their own concerns — yet the epidemiological ledger is not close. Waterborne typhoid and cholera killed thousands of Britons every year before treatment; they kill effectively none now. A good answer states a risk, states the benefit, and commits to the comparison rather than hiding behind 'there are pros and cons'.
Prove the disproportionation, the way a mark scheme wants it. In Cl₂, chlorine is 0. In HClO: hydrogen +1 and oxygen −2, so chlorine must be +1 — oxidised, 0 → +1. In HCl: chlorine is −1 — reduced, 0 → −1. Both changes happen to the same element in one reaction, which is the definition of disproportionation (and quoting both state changes with numbers is what earns the mark). Contrast the sunlight reaction: there chlorine only falls, 0 → −1, while oxygen rises, −2 → 0 — a redox reaction, but not a disproportionation, because the two changes happen to different elements. Being able to sort those two reactions is a favourite discriminator question.
CaseRequired practical 4 — the order of the tests is the method
RP4 asks you to identify unknown ions with nothing but test tubes: Group 2 cations and ammonium on one side, carbonate, sulfate and halides on the other. The cation tests recycle 3.2.2's trends: sodium hydroxide gives a thick white precipitate for magnesium (insoluble in excess), fading to little or none down the group as the hydroxides turn soluble, while dilute sulfuric acid runs the opposite ladder — nothing for magnesium, dense white for barium. Each metal returns a unique pair of results. Ammonium is the odd one out: warm the solution with sodium hydroxide and test the gas — ammonia turns damp red litmus paper blue, the only common gas that does.
The anion tests you now know individually; RP4's real content is their sequence. Test for carbonate first: add dilute hydrochloric acid and look for effervescence, confirming the gas as CO₂ with limewater. Then sulfate: acidified barium chloride, white precipitate. Then halides last: acidified silver nitrate, followed by the ammonia ladder. The order is not administrative — it is the chemistry. Carbonate would fake a positive in both later tests (barium carbonate and silver carbonate are both insoluble), and sulfate would fake the silver test (silver sulfate precipitates too). Each test therefore eliminates an ion that could forge the results of the next, and 'explain why the tests are done in this order' is the six-mark question this practical exists to set up.
Technique notes that earn method marks: use a fresh portion of solution for each test, add reagents dropwise, use deionised water throughout (tap water carries enough chloride to fog a silver nitrate test), and record what you see — 'no visible change' is data; a blank row is not.
VocabularyKey terms the mark scheme pays for
TrapsMisconceptions that cost marks
ExamWhat examiners want
This section is scored in observations, and AQA wants the language of eyes, not inference: 'misty fumes', 'brown gas', 'violet vapour', 'pale yellow solid', 'smell of bad eggs', 'white precipitate, insoluble in dilute but soluble in concentrated ammonia'. Writing 'HBr gas is produced' where the question asked what you would SEE routinely loses the mark. Build the sulfuric acid ladder as a table in your head — halide × observations × sulfur product — because six-markers ask you to compare all three halides, and the marks land one observation and one oxidation state at a time.
Trend explanations run the standard machine — shells, shielding, distance, attraction — but check the direction of the question first: halogens as oxidisers weaken down the group, halide ions as reducers strengthen. For disproportionation, quote both oxidation-state changes with numbers (0 → +1 and 0 → −1); the word alone, without the bookkeeping, is not a complete answer.
On the practical side, every 'why' has a one-line answer you should pre-load: why nitric acid before silver nitrate (destroys carbonate/hydroxide without adding halide); why hydrochloric before barium chloride (removes carbonate/sulfite without adding sulfate); why carbonate is tested first in RP4 (it fakes both later precipitates). And the water-treatment evaluation wants a committed judgement: name a real risk (toxic gas, chlorinated by-products), name the benefit (elimination of waterborne typhoid and cholera), and conclude that the benefit dominates — examiners reward the decision, not the fence-sit.