S4 Chemistry Practice Paper: Redox, Chemical Cells and Electrolysis (100 marks)

2026-11-10

S4 Chemistry · Redox and electrochemistry

S4 Chemistry paper: 6 questions, 26 parts, 100 marks, 75 min. Covers oxidation numbers, half-equations, chemical cells, electrolysis, Faraday and mole calculations.

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Section A — Oxidation numbers and redox

Show your working and give the sign of each oxidation number.

#	Question	Answer and marking notes
1	Find the oxidation number of manganese in KMnO4.	Let x be the oxidation number of Mn: +1 + x + 4 × (−2) = 0, so x = +7.
2	Find the oxidation number of chromium in K2Cr2O7.	Let x be the oxidation number of Cr: 2 × (+1) + 2x + 7 × (−2) = 0, so x = +6.
3	Is the change Fe2+ → Fe3+ an oxidation or a reduction? Write the half-equation, including the electron.	The oxidation number of iron rises from +2 to +3, so electrons are lost: it is an oxidation. The half-equation is Fe2+ → Fe3+ + e−.
4	In the reaction Zn + Cu2+ → Zn2+ + Cu, which species is the oxidising agent?	Cu2+ is the oxidising agent; it gains the electrons released by zinc and is reduced to copper.

Section B — Half-equations and ionic equations

Assume acidic solution and state the number of electrons clearly.

#	Question	Answer and marking notes
5	Write the half-equation for the reduction of MnO4− to Mn2+ in acidic solution, showing the electrons.	MnO4− + 8H+ + 5e− → Mn2+ + 4H2O (Mn falls from +7 to +2, so 5 electrons are gained).
6	Write the full ionic equation between Fe2+ and MnO4− in acidic solution, and state their mole ratio.	5Fe2+ + MnO4− + 8H+ → 5Fe3+ + Mn2+ + 4H2O; the mole ratio of Fe2+ to MnO4− is 5 : 1.
7	How many moles of electrons are accepted when 1 mol of MnO4− is reduced completely to Mn2+?	Each mole of MnO4− accepts 5 mol of electrons.
8	Which species in the reaction above is the reducing agent? Explain.	Fe2+ is the reducing agent: it donates electrons to MnO4− and is itself oxidised to Fe3+.

Section C — Chemical cells

Name the electrodes and state the direction of electron flow.

#	Question	Answer and marking notes
9	When a zinc-copper cell discharges, write the reaction at each electrode.	Zinc electrode: Zn → Zn2+ + 2e− (oxidation); copper electrode: Cu2+ + 2e− → Cu (reduction).
10	In the zinc-copper cell, which electrode is the negative terminal and which is the positive terminal?	Zinc is the negative terminal (anode, where oxidation occurs); copper is the positive terminal (cathode, where reduction occurs).
11	In which direction do electrons flow through the external circuit?	Electrons flow from the zinc electrode through the external circuit to the copper electrode.
12	What is the function of the salt bridge in a cell?	The salt bridge completes the internal circuit and allows ions to migrate so that both solutions stay electrically neutral.
13	The standard electrode potential of Zn2+/Zn is -0.76 V and that of Cu2+/Cu is +0.34 V. Find the electromotive force of the cell.	Electromotive force = positive − negative = 0.34 − (-0.76) = 1.1 V.

Section D — Electrolysis and charge calculations

Take the Faraday constant as 96 500 C/mol and Cu = 63.5.

#	Question	Answer and marking notes
14	Copper electrodes are used to electrolyse copper(II) sulfate solution. Write the anode reaction.	Anode: Cu → Cu2+ + 2e−; the copper electrode itself dissolves.
15	Write the cathode reaction for the same electrolysis.	Cathode: Cu2+ + 2e− → Cu; copper is deposited on the cathode.
16	A current of 2 A passes for 30 minutes. Find the quantity of charge that flows.	Q = It = 2 × 1800 = 3600 C.
17	Find the mass of copper deposited at the cathode (3 significant figures).	Moles of electrons = 3600 ÷ 96 500 = 0.03731 mol; moles of copper = 0.03731 ÷ 2 = 0.01865 mol; mass = 0.01865 × 63.5 = 1.18 g.
18	Why does the blue colour of the solution stay roughly the same when copper electrodes are used?	The amount of copper dissolving at the anode equals the amount deposited at the cathode, so the concentration of copper ions, and hence the blue colour, stays roughly the same.

Section E — Applications of redox reactions

Give the relevant half-equation or principle.

#	Question	Answer and marking notes
19	Which two conditions are needed for iron to rust?	Both oxygen (air) and water are needed.
20	Write the half-equation for the oxidation of iron to iron(II) ions.	Fe → Fe2+ + 2e−.
21	Why does coating iron with zinc prevent rusting?	Zinc is more reactive than iron and is oxidised in preference to it, so it protects the iron; even when the coating is scratched, zinc acts as a sacrificial anode.
22	On what principle does bleach disinfect?	The chlorate(I) ion in bleach is a strong oxidising agent that oxidises bacterial proteins and enzymes, destroying their activity.

Section F — Mole and concentration calculations

State units and give masses to 3 significant figures.

#	Question	Answer and marking notes
23	0.1 mol of KMnO4 is reduced completely to Mn2+. How many moles of electrons are accepted?	Moles of electrons = 0.1 × 5 = 0.50 mol.
24	How many moles of KMnO4 are there in 25.0 cm3 of 0.02 mol/dm3 KMnO4 solution?	Volume = 25.0 cm3 = 0.025 dm3; moles = 0.025 × 0.02 = 0.00050 mol.
25	What volume of 0.1 mol/dm3 Fe2+ solution can this amount of KMnO4 oxidise?	Moles of Fe2+ = 0.00050 × 5 = 0.00250 mol; volume = 0.00250 ÷ 0.1 = 0.025 dm3 = 25.0 cm3.
26	What mass of iron is present in 0.00250 mol of Fe2+ (Fe = 55.8)?	Mass = 0.00250 × 55.8 = 0.1395 g ≈ 0.140 g.

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