HookCamelford, 1988: twenty tonnes of aluminium sulfate in the wrong tank

On 6 July 1988, a relief tanker driver arrived at the Lowermoor water treatment works in North Cornwall, found the site unmanned, and emptied his load — twenty tonnes of concentrated aluminium sulfate solution — into what he took to be the storage tank. It was the tank feeding treated water to around 20,000 people in and around Camelford. Within hours residents were reporting water that tasted foul and curdled the milk in their tea; the acidity stripped copper from household pipes, and some unlucky blondes found their hair tinged green with dissolved copper. It remains Britain's worst mass water-contamination incident.

The chemistry of that week is the chemistry of this section. Dissolve an aluminium salt in water and you do not simply get neutral 'aluminium solution' — you get the hexaaqua ion \([\text{Al}(\text{H}_2\text{O})_6]^{3+}\), and that ion is a genuine acid, donating protons from its own water ligands until the tap water at Camelford ran towards pH 4. Why 3+ aqua ions are so much more acidic than 2+ ones, what hydroxide, ammonia and carbonate do to each of them, and how those reactions become a test-tube identification scheme (Required Practical 11) — that is 3.2.6. It is the section where the colours finally earn their keep: every precipitate is a clue.

ModelMetal-aqua ions are acids — and charge density sets how strong

In aqueous solution the ions this section runs on are six-coordinate aqua complexes: green \([\text{Fe}(\text{H}_2\text{O})_6]^{2+}\), blue \([\text{Cu}(\text{H}_2\text{O})_6]^{2+}\), colourless \([\text{Al}(\text{H}_2\text{O})_6]^{3+}\), and \([\text{Fe}(\text{H}_2\text{O})_6]^{3+}\) — pale violet in theory, though real iron(III) solutions look yellow-brown because hydrolysis has already begun.

Hydrolysis is the key reaction: a water ligand surrenders a proton to a solvent water molecule, \([\text{Fe}(\text{H}_2\text{O})_6]^{3+} + \text{H}_2\text{O} \rightleftharpoons [\text{Fe}(\text{H}_2\text{O})_5(\text{OH})]^{2+} + \text{H}_3\text{O}^{+}\). Note the reversible arrows — this is an equilibrium, lying only partly to the right, which is why these solutions are weakly acidic rather than hydrochloric-acid acidic. Solutions of iron(III) salts sit around pH 2–3; iron(II) and copper(II) salts nearer pH 5–6.

The difference is charge density — the charge-to-size ratio of the central ion. A 3+ ion is more highly charged and (for these metals) smaller than a 2+ ion, so it polarises its water ligands more strongly: electron density is pulled from the O–H bonds, the bonds weaken, and the proton leaves more readily. Higher charge density, stronger acid. That single argument is the AO1 backbone of the whole section — and it is exactly why Camelford's aluminium sulfate, a 3+ salt delivered twenty tonnes at a time, sent the water supply acidic.

MechanismHydroxide and ammonia — the same precipitates, until you add excess

Add sodium hydroxide solution dropwise to any of the four aqua ions and you see the same type of reaction: hydroxide ions strip protons from the water ligands, stepwise, until the complex is electrically neutral — and a neutral complex is insoluble, so it precipitates. For copper: \([\text{Cu}(\text{H}_2\text{O})_6]^{2+} + 2\text{OH}^- \rightarrow \text{Cu}(\text{H}_2\text{O})_4(\text{OH})_2 + 2\text{H}_2\text{O}\), a pale blue precipitate. Iron(II) gives green \(\text{Fe}(\text{H}_2\text{O})_4(\text{OH})_2\), which darkens at the surface on standing as air oxidises it towards brown iron(III) hydroxide. Iron(III) gives brown \(\text{Fe}(\text{H}_2\text{O})_3(\text{OH})_3\); aluminium gives white \(\text{Al}(\text{H}_2\text{O})_3(\text{OH})_3\).

Ammonia, added dropwise, is also just a base here — \([\text{Fe}(\text{H}_2\text{O})_6]^{2+} + 2\text{NH}_3 \rightarrow \text{Fe}(\text{H}_2\text{O})_4(\text{OH})_2 + 2\text{NH}_4^{+}\) — so the initial precipitates are identical to sodium hydroxide's. The diagnosis lives in the excess.

Excess sodium hydroxide dissolves only the aluminium precipitate: \(\text{Al}(\text{H}_2\text{O})_3(\text{OH})_3 + \text{OH}^- \rightarrow [\text{Al}(\text{OH})_4]^- + 3\text{H}_2\text{O}\), a colourless solution — aluminium hydroxide is amphoteric, acting as an acid towards the extra hydroxide (and as a base towards acids, re-dissolving in dilute HCl). Excess ammonia dissolves only the copper precipitate, and by a different mechanism entirely — ligand substitution: \(\text{Cu}(\text{H}_2\text{O})_4(\text{OH})_2 + 4\text{NH}_3 \rightarrow [\text{Cu}(\text{NH}_3)_4(\text{H}_2\text{O})_2]^{2+} + 2\text{OH}^- + 2\text{H}_2\text{O}\), the deep royal-blue solution. One scheme, two 'dissolves in excess' outcomes, two different kinds of chemistry — that distinction is a favourite six-mark discriminator.

MechanismCarbonate — the fizz that tells you the charge

Sodium carbonate splits the four ions cleanly in two, and the split is the acidity argument made visible.

The 2+ aqua ions are too weakly acidic to attack carbonate, so the carbonate ion simply partners the metal ion and precipitates: \([\text{Cu}(\text{H}_2\text{O})_6]^{2+} + \text{CO}_3^{2-} \rightarrow \text{CuCO}_3 + 6\text{H}_2\text{O}\) (blue-green), and iron(II) gives green FeCO\(_3\) the same way. No gas, no drama — just a carbonate precipitate.

The 3+ aqua ions are strong enough acids to destroy carbonate, protonating it to carbon dioxide exactly as vinegar attacks baking soda: \(2[\text{Al}(\text{H}_2\text{O})_6]^{3+} + 3\text{CO}_3^{2-} \rightarrow 2\text{Al}(\text{H}_2\text{O})_3(\text{OH})_3 + 3\text{CO}_2 + 3\text{H}_2\text{O}\) — a hydroxide precipitate (white here, brown for iron(III)) plus effervescence. This is why aluminium carbonate and iron(III) carbonate do not exist as stable compounds: the 3+ aqua ion's acidity dismantles the carbonate on contact. In the exam, 'bubbles with sodium carbonate solution' translates directly to 'the ion is 3+' — the single most decisive observation in the identification scheme.

DataRequired Practical 11 — running the identification scheme properly

RP11 turns the last two blocks into a method. Take a fresh portion of the unknown solution for every test — a precipitate from one reagent contaminates the next observation. Note the solution's own colour first (blue suggests copper, green iron(II), yellow-brown iron(III), colourless aluminium — or nothing transition at all). Then run three tests: sodium hydroxide dropwise until a precipitate forms, then to excess; ammonia dropwise, then to excess; sodium carbonate solution to a fresh sample, watching for effervescence.

Technique is where marks (and real identifications) are lost. Add reagents dropwise: pour in excess at once and a precipitate that dissolves in excess — aluminium's with NaOH, copper's with NH\(_3\) — may never visibly form, and you record a false negative. Record observations immediately: iron(II)'s green hydroxide begins darkening within minutes as air oxidises it, and a late reading looks like iron(III). Use colour words the scheme uses — 'white precipitate, insoluble in excess ammonia' — not impressions ('cloudy'). And state what does NOT happen: 'no effervescence with carbonate' is the evidence that rules out a 3+ ion.

Worked example

Solution X is pale green. With NaOH it gives a green precipitate, insoluble in excess, which turns brown at its surface after standing in air. With sodium carbonate it gives a green precipitate and no effervescence. Solution Y is colourless. With NaOH it gives a white precipitate that dissolves in excess to a colourless solution; with sodium carbonate it gives a white precipitate with vigorous effervescence. Identify X and Y with equations.

X: the green solution and green hydroxide point to iron(II); insolubility in excess rules out aluminium; no fizz with carbonate confirms a 2+ ion, and the surface browning is Fe(II) oxidising to Fe(III) in air. X contains \([\text{Fe}(\text{H}_2\text{O})_6]^{2+}\): \([\text{Fe}(\text{H}_2\text{O})_6]^{2+} + 2\text{OH}^- \rightarrow \text{Fe}(\text{H}_2\text{O})_4(\text{OH})_2 + 2\text{H}_2\text{O}\), and with carbonate \([\text{Fe}(\text{H}_2\text{O})_6]^{2+} + \text{CO}_3^{2-} \rightarrow \text{FeCO}_3 + 6\text{H}_2\text{O}\).

Y: a white precipitate dissolving in excess NaOH means an amphoteric hydroxide, and effervescence with carbonate means a 3+ aqua ion — together, aluminium. Y contains \([\text{Al}(\text{H}_2\text{O})_6]^{3+}\): \(\text{Al}(\text{H}_2\text{O})_3(\text{OH})_3 + \text{OH}^- \rightarrow [\text{Al}(\text{OH})_4]^- + 3\text{H}_2\text{O}\) for the excess step, and \(2[\text{Al}(\text{H}_2\text{O})_6]^{3+} + 3\text{CO}_3^{2-} \rightarrow 2\text{Al}(\text{H}_2\text{O})_3(\text{OH})_3 + 3\text{CO}_2 + 3\text{H}_2\text{O}\) for the fizz. Chain every observation to an inference — that reasoning structure, not the bare names, is what the 5- and 6-mark versions of this question pay for.

VocabularyKey terms the mark scheme pays for

Hexaaqua ion

A metal ion coordinated by six water ligands, e.g. [Fe(H₂O)₆]²⁺ (green), [Cu(H₂O)₆]²⁺ (blue), [Fe(H₂O)₆]³⁺ (violet, appearing yellow-brown), [Al(H₂O)₆]³⁺ (colourless).

Hydrolysis (acidity) reaction

A water ligand donating a proton to solvent water, e.g. [Fe(H₂O)₆]³⁺ ⇌ [Fe(H₂O)₅(OH)]²⁺ + H₃O⁺ — an equilibrium, so the solutions are weakly acidic.

Charge density

An ion's charge relative to its size. High charge density (3+, small radius) polarises water ligands, weakens their O–H bonds and makes the aqua ion a stronger acid.

Neutral hydroxide complex

The uncharged species formed when enough protons are stripped away, e.g. Cu(H₂O)₄(OH)₂ or Al(H₂O)₃(OH)₃ — no overall charge, so it is insoluble and precipitates.

Amphoteric hydroxide

A hydroxide that reacts with both bases and acids: Al(H₂O)₃(OH)₃ dissolves in excess NaOH as [Al(OH)₄]⁻ and in dilute acid back to [Al(H₂O)₆]³⁺.

Ligand substitution (in excess ammonia)

The one step in this scheme that is not acid-base: excess NH₃ replaces ligands on copper to give deep blue [Cu(NH₃)₄(H₂O)₂]²⁺, dissolving the pale blue precipitate.

Effervescence with carbonate

The 3+ diagnostic: aqua ions of Al³⁺ and Fe³⁺ are acidic enough to protonate CO₃²⁻ to CO₂ gas, precipitating the hydroxide — 2+ ions just precipitate the metal carbonate.

TrapsMisconceptions that cost marks

“Sodium hydroxide and ammonia give different precipitates, so they test for different things.”

Actually: Added dropwise they are the same test — both act as bases and produce the identical hydroxide precipitate. All the diagnostic power is in the EXCESS step: excess NaOH redissolves only aluminium's precipitate, excess NH₃ only copper's.

“Iron(III) carbonate forms first and then decomposes to give the fizz.”

Actually: It never forms. The [Fe(H₂O)₆]³⁺ ion is acidic enough to protonate carbonate ions on contact, so the products are the hydroxide precipitate and CO₂ directly — which is precisely why no stable 3+ metal carbonates exist.

“Aluminium's precipitate dissolving in excess NaOH is ligand substitution, like copper in ammonia.”

Actually: Different chemistry. Aluminium hydroxide dissolving is a further acid-base (deprotonation) step — Al(H₂O)₃(OH)₃ acting as an acid to give [Al(OH)₄]⁻. Copper in excess ammonia is genuine ligand substitution, with NH₃ replacing ligands around the metal. Naming the correct reaction type is where the mark sits.

ExamWhat examiners want

Observations are marked as three-part sentences: colour + species + behaviour in excess — 'pale blue solution gives a pale blue precipitate, insoluble in excess sodium hydroxide'. AQA's examiner reports repeatedly warn that 'cloudy' and 'goes darker' earn nothing; the scheme's own colour words do. When a question asks for an equation, write the full complex formulas with charges and check the charge balance line by line — losing a 2+ from one side of a deprotonation equation is the commonest slip in the topic.

For 'explain why the 3+ solution is more acidic' questions, the creditable chain is: higher charge and smaller radius → higher charge density → greater polarisation of coordinated water → O–H bonds weakened → protons released more readily. Write hydrolysis equations with ⇌, because stating that the equilibrium lies well to the left is itself a marking point — these are weak acids, not strong ones.

In RP11-style questions, the practical marks are for fresh samples for each test, dropwise addition before excess, immediate recording (iron(II)'s precipitate oxidises brown on standing), and negative observations stated explicitly. In deduction questions, structure the answer as observation → inference for every line before naming the ion: the reasoning chain scores even where the final identification slips, and a bare name without reasoning scores almost nothing.