KCSE 2025 Chemistry P2 Q4 — Chlorine: Preparation and Reactions
Published
The Question
“(a) Describe how chlorine gas is prepared in the laboratory, starting from solid sodium chloride and concentrated sulphuric acid and passing the gas produced over heated manganese(IV) oxide; write the equation for the oxidation step. (b) A flow chart shows three reactions of chlorine: to the right it reacts with iron to form compound C; downward, in step II, it forms sodium chlorate; to the left it reacts with phosphorus to form compounds D and E. (i) Identify compound C. (ii) Name the reagent and conditions in step II and identify the product. (iii) Identify D and E. (c) Write the equation for the reaction of chlorine with aqueous hydrogen sulphide. (d) State the observation when chlorine is (i) bubbled into potassium iodide solution and (ii) placed in contact with moist red litmus paper. (e) Write the ionic equation for the reaction when lead(II) nitrate solution is added to potassium chloride solution. (f) Give one use of chlorine.”
(a) Preparing chlorine in the laboratory
Chlorine is made in two stages. First, concentrated sulphuric acid is added to solid sodium chloride and the mixture is heated, giving off hydrogen chloride gas. The hydrogen chloride is then passed over heated manganese(IV) oxide, which oxidises it to chlorine gas, also forming manganese(II) chloride and water. The oxidation step is the key equation.
(b)(i) Reaction with iron — compound C
Chlorine is a strong oxidising agent, so when it passes over hot iron it forces the iron up to its highest oxidation state, iron(III). Compound C is iron(III) chloride, FeCl₃.
(b)(ii) Step II — reaction with hot concentrated alkali
The reagent in step II is sodium hydroxide, and the condition is decisive: it must be hot and concentrated. Chlorine with cold dilute alkali gives the chloride and chlorate(I), but with hot concentrated sodium hydroxide it gives sodium chlorate(V), NaClO₃, along with sodium chloride and water.
(b)(iii) Reaction with phosphorus — compounds D and E
Chlorine reacts with phosphorus to give two chlorides depending on how much chlorine is present. With a limited supply of chlorine the product is phosphorus trichloride, PCl₃; with excess chlorine it is phosphorus pentachloride, PCl₅. So D and E are PCl₃ and PCl₅.
(c) Reaction with hydrogen sulphide
Chlorine acts as the oxidising agent again: bubbled into aqueous hydrogen sulphide, it takes the hydrogen and sets the sulphur free as a pale yellow solid, while forming hydrochloric acid.
(d)(i) Observation with potassium iodide solution
Chlorine is more reactive than iodine, so when bubbled into colourless potassium iodide solution it displaces the iodine. The observation is that the colourless solution turns brown (yellow-brown) as iodine is set free.
(d)(ii) Observation with moist red litmus paper
You might expect red litmus to stay red in an acidic gas, but moist chlorine forms chloric(I) acid, HOCl, which is a bleach. The observation is that the moist red litmus paper is bleached, turning white.
(e) Ionic equation for the lead(II) chloride precipitate
When lead(II) nitrate solution is added to potassium chloride solution, a white precipitate of lead(II) chloride forms. In the ionic equation only the ions that actually react are written — the lead(II) ions and the chloride ions — leaving out the spectator potassium and nitrate ions.
(f) A use of chlorine
Chlorine is added to drinking water in treatment works because it kills harmful bacteria and micro-organisms, sterilising the water and making it safe to drink.
Final Result
(a) Warm concentrated sulphuric acid with solid sodium chloride to make hydrogen chloride, then pass the HCl over heated MnO₂: 4HCl + MnO₂ → MnCl₂ + 2H₂O + Cl₂. (b) With iron, compound C is FeCl₃; step II uses hot concentrated NaOH to give sodium chlorate(V), NaClO₃; with phosphorus, D and E are PCl₃ and PCl₅. (c) Cl₂ + H₂S → S + 2HCl. (d) With KI(aq) the colourless solution turns brown as iodine is displaced; moist red litmus is bleached white. (e) Pb²⁺(aq) + 2Cl⁻(aq) → PbCl₂(s). (f) Chlorine kills bacteria and micro-organisms in the treatment of drinking water.
Why this method works
The whole question turns on chlorine being a strong oxidising agent and a reactive non-metal. Its laboratory preparation is itself a redox step: manganese(IV) oxide oxidises the chloride in hydrogen chloride up to chlorine gas. That same oxidising power drives most of its reactions — it pushes iron to its highest state, iron(III); it strips hydrogen from hydrogen sulphide to release sulphur; and it displaces the less reactive iodine from iodide because chlorine sits higher in the halogen reactivity series. The alkali and phosphorus reactions show how conditions decide the product: hot concentrated alkali gives chlorate(V) where cold dilute alkali gives chlorate(I), and the amount of chlorine decides PCl₃ versus PCl₅. The litmus result is a classic trap — moist chlorine bleaches rather than simply reddening, because it forms chloric(I) acid, and that very reactivity toward micro-organisms is exactly why chlorine sterilises drinking water.
Each equation balances: 4HCl + MnO₂ gives 4 H, 4 Cl, 1 Mn and 2 O on both sides. In the displacement, chlorine lying above iodine in the halogen group guarantees I₂ is set free. In the ionic equation the charges balance — 2+ against two 1− ions gives zero — matching the neutral solid PbCl₂.