GCSE Combined Science Acids, electrolysis & salts revision

Solutions run from strongly acidic (pH 0) through neutral (pH 7 — pure water) to strongly alkaline (pH 14). The number reports the concentration of hydrogen ions, H⁺(aq): acids release H⁺ into water, and the more they release, the lower the pH. Alkaline solutions are rich in hydroxide ions, OH⁻(aq). Universal indicator steps through red → green → purple as pH climbs, while a pH probe reads the number directly and more precisely.

Mix an acid with an alkali and the H⁺ and OH⁻ ions pair off into water: H⁺(aq) + OH⁻(aq) → H₂O(l) — neutralisation. The leftover ions (say Na⁺ and Cl⁻) stay dissolved as a salt, which is why the full story is *acid + alkali → salt + water*. Farmers use the same chemistry on a field scale — calcium hydroxide (liming) lifts the pH of acidic soil — and antacid tablets use mild bases to mop up excess stomach acid.

(Higher) Strong acids — hydrochloric, sulfuric, nitric — ionise completely in water: every molecule releases its H⁺. Weak acids such as ethanoic (vinegar) or citric ionise only partially, so at the same concentration they hold back most of their H⁺. Strong and weak describe *ionisation*; concentrated and dilute describe *how much acid is dissolved* — a dilute strong acid and a concentrated weak acid can share a pH.

Every acid reaction follows a template — learn the template and the products write themselves:

Naming the salt is a two-part job: take the metal from the base, carbonate or metal, and the family from the acid. Hydrochloric acid always makes chlorides, sulfuric acid makes sulfates, nitric acid makes nitrates. So magnesium + hydrochloric acid can only be magnesium chloride, no matter how the question is dressed up.

Trilogy’s salt practical starts from an insoluble base — copper(II) oxide is the classic — so every leftover is removable by filtration. The five steps:

1. 11. Warm the dilute acid

   Gently heat about 25 cm³ of dilute sulfuric acid in a beaker. Warm — never boil — so the reaction speeds up without the acid spitting or evaporating away.

2. 22. Add the base in excess

   Stir in black copper(II) oxide a spatula at a time; it dissolves as it reacts and the solution deepens to blue. Stop when some black solid refuses to dissolve — that proves every drop of acid has been used up.

3. 33. Filter

   Filter the warm mixture. The filtrate is blue copper sulfate solution; the residue is the excess oxide. Filtering removes the impurity you deliberately created.

4. 44. Evaporate to crystallisation

   Heat the filtrate until roughly a third of the water has gone and crystals just start forming at the edges. Do not evaporate to dryness.

5. 55. Cool and dry

   Leave the solution to cool slowly so large, pure crystals grow, then pat them dry between filter papers.

The independent variable in an investigation of this practical might be the metal oxide used; control variables include the acid volume and concentration, the temperature and the stirring rate. Hazards: hot acid and hot glassware (goggles, keep the acid below boiling), and copper salts are harmful to aquatic life — small spills to the sink, not the bin, and never down a storm drain.

Electrolysis is decomposition by electricity. A solid ionic compound is useless — its ions are locked in the lattice — so the compound must be molten or dissolved for the ions to migrate. The cathode is the negative electrode and attracts positive ions (cations); the anode is positive and attracts negative ions (anions) — remember PANIC: Positive Anode, Negative Is Cathode.

For a molten binary ionic compound, prediction is easy: the metal appears at the cathode and the non-metal at the anode. Molten lead bromide gives silvery lead beads at the cathode and brown bromine vapour at the anode; molten sodium chloride gives sodium and green chlorine.

Extracting aluminium shows electrolysis earning its (high) energy bill. Aluminium sits above carbon, so carbon reduction cannot touch it — its ore, aluminium oxide (alumina), must be melted and electrolysed. Pure alumina melts at about 2072 °C, far too hot to run a cell, so cryolite is added and the cell runs near 950 °C. Aluminium forms at the cathode; oxygen forms at the anode, where it attacks the hot carbon anodes, burning them away so they must be replaced regularly. The cost drivers are the heat and the huge continuous electric current — which is why aluminium smelters sit next to cheap power.

In solution there is always water competing for the electrodes, so two rules decide the products with inert electrodes: at the cathode, hydrogen forms unless the metal is less reactive than hydrogen (copper, silver — those plate out instead); at the anode, a halide (chloride, bromide, iodide) gives the halogen if present, otherwise oxygen forms from hydroxide ions.

Apply the rules to the Trilogy set: concentrated sodium chloride (brine) gives hydrogen at the cathode, chlorine at the anode, and sodium hydroxide left in solution — three useful products at once, and chlorine’s pale green gas bleaches damp litmus. Copper sulfate plates pink-brown copper on the cathode (the blue fades as Cu²⁺ is used up) and bubbles oxygen at the anode. Dilute sulfuric acid is essentially electrolysis of water: hydrogen and oxygen come off in a 2 : 1 volume ratio, matching water’s two-to-one hydrogen-to-oxygen atoms.

(Higher) Every electrode reaction can be written as a half equation that balances atoms and charge. At the cathode, positive ions gain electrons — reduction: 2H⁺ + 2e⁻ → H₂, Cu²⁺ + 2e⁻ → Cu, Al³⁺ + 3e⁻ → Al. At the anode, negative ions lose electrons — oxidation: 2Cl⁻ → Cl₂ + 2e⁻, 2Br⁻ → Br₂ + 2e⁻, and for oxygen from hydroxide ions: 4OH⁻ → O₂ + 2H₂O + 4e⁻. Check every half equation the same way: atoms equal on both sides, and the charges must cancel too.

The electron language also explains the observations. Dilute sulfuric acid gives twice the volume of hydrogen as oxygen because four electrons buy one O₂ but two H₂ molecules. The blue of copper sulfate solution fades because Cu²⁺ ions are removed at the cathode faster than anything replaces them. Chlorine’s bleaching power comes from the reactive Cl₂ molecules made when chloride ions surrender their electrons.