OxfordAQA International GCSE Chemistry: Factors affecting equilibrium, and production of ammonia and sulfuric acid (9202) – Practice Questions

These are original questions written for Marlbridge, for revision and practice on this content. They are not reproduced past-paper questions, and they do not replicate the exam’s exact structure, question count or mark tariffs – examination boards hold copyright in their own papers. Use these alongside the official past papers from your board or school.

Written for OxfordAQA International GCSE Chemistry (9202), this set follows the OxfordAQA International GCSE Chemistry (9202) specification, Version 6.2, for exams May/June 2018 onwards. It covers section 3.8.2 Factors affecting equilibrium (statements a to e) and section 3.8.3 Production of ammonia and sulfuric acid (statements a to d). Neither section has a required practical. The qualification is untiered, so every question suits every candidate. Use relative atomic masses H = 1, N = 14, O = 16 and S = 32 where needed.

Links: study guide · revision notes · course hub · checklist · free diagnostics

Questions

1. A reversible reaction is carried out in a sealed flask.

(a) State what is meant by equilibrium for this reaction. [2] (b) A student writes: “At equilibrium the reaction has finished and there are equal amounts of reactants and products.” Identify the two mistakes in this statement. [2]

2. Carbon monoxide and chlorine combine to form carbonyl chloride, COCl₂:

CO(g) + Cl₂(g) ⇌ COCl₂(g), with energy given out in the forward direction.

(a) What happens to the yield of COCl₂ if the temperature is raised? Explain your answer. [2] (b) What happens to the yield of COCl₂ if the pressure is increased? Explain your answer. [2]

3. Nitrosyl chloride decomposes when heated:

2NOCl(g) ⇌ 2NO(g) + Cl₂(g). The forward reaction is endothermic.

(a) State whether a high or a low temperature gives the larger yield of chlorine. Explain your answer. [2] (b) State whether a high or a low pressure gives the larger yield of chlorine. Explain your answer. [2] (c) For the reaction 2HBr(g) ⇌ H₂(g) + Br₂(g), explain why changing the pressure does not change the yield of hydrogen. [1]

4. A gas reaction is G(g) + 2J(g) ⇌ K(g). The table shows the percentage of K in the equilibrium mixture under four sets of conditions.

Temperature	% K at 1 atmosphere	% K at 10 atmospheres
100 °C	52	71
300 °C	30	48

(a) Is the forward reaction exothermic or endothermic? Use the table to decide, and explain how. [2] (b) Explain how the data support the rule about pressure and numbers of molecules. [2] (c) Suggest why a factory making K might not choose 100 °C, even though it gives a higher yield. [1]

5. Describe how ammonia is manufactured by the Haber process. Include the raw materials and their sources, the conditions, and what happens to the gases after they leave the reactor. [6]

6. The forward reaction in the Haber process is exothermic.

(a) Suggest why the plant runs at about 450 °C and not at a cooler temperature. [3] (b) Give reasons for running at about 200 atmospheres rather than at a much higher pressure. [3]

7. Most ammonia is used to make fertilizers.

(a) Explain why there is a global need for fertilizers. [3] (b) Ammonium sulfate is made by the reaction 2NH₃ + H₂SO₄ → (NH₄)₂SO₄. Calculate the mass of ammonia needed to react with 49 tonnes of sulfuric acid, and the mass of ammonium sulfate formed. [3]

8. Sulfuric acid is made by the contact process.

(a) Give balanced equations, including state symbols, for Stages 1 and 3. [2] (b) State the catalyst, the temperature and the pressure used in Stage 2. [3]

9. Stage 2 of the contact process is exothermic in the forward direction (equation: 2SO₂(g) + O₂(g) ⇌ 2SO₃(g)). Explain why a catalyst, a temperature of about 450 °C and atmospheric pressure are used. [6]

10. In one pass through a Haber reactor, 560 kg of nitrogen enters and 15% of it reacts.

(a) Calculate the mass of nitrogen that reacts. [1] (b) Calculate the mass of ammonia formed. [2] (c) Calculate the mass of hydrogen that reacts. [1] (d) State what happens to the nitrogen that does not react, and explain why this matters. [2]

11. A contact process plant starts with 4.8 tonnes of sulfur.

(a) Calculate the mass of sulfur dioxide formed in Stage 1. [2] (b) Calculate the mass of oxygen needed in Stage 2 to convert all of this sulfur dioxide to sulfur trioxide. [2] (c) Calculate the maximum mass of sulfuric acid that can be made. [2] (d) Suggest one reason why the plant makes less than this maximum mass. [1]

Answers

1. (a) The forward and backward reactions are both happening [1] at exactly the same rate (in a closed system) [1]. (b) The reaction has not finished: both reactions are still taking place [1]. The amounts are constant, not necessarily equal [1]. Examiner insight: “The rates are equal” alone is often judged incomplete; also say that both reactions are still going on.

2. (a) The yield of COCl₂ decreases [1]. The forward reaction is exothermic, so the backward reaction is endothermic, and raising the temperature increases the yield from the endothermic reaction [1]. (b) The yield of COCl₂ increases [1]. There are 2 molecules on the left and 1 on the right, and increased pressure favours the reaction producing fewer molecules [1]. Examiner insight: Tie each explanation to this equation (the endothermic direction, or the molecule count); “it shifts” alone is unlikely to gain credit.

3. (a) High temperature [1]. The forward reaction is endothermic, and raising the temperature increases the yield from the endothermic reaction [1]. (b) Low pressure [1]. 2 molecules become 3, and decreased pressure favours the reaction producing the greater number of molecules [1]. (c) There are 2 molecules of gas on each side, so neither direction produces fewer molecules [1]. Examiner insight: Show the molecule counts from the equation, such as “2 → 3”; a bare “low pressure” with no count rarely earns the explanation mark.

4. (a) At each pressure, % K falls as the temperature rises [1]. Raising the temperature favours the endothermic reaction, so the backward reaction is endothermic and the forward reaction is exothermic [1]. (b) At each temperature, % K is higher at 10 atmospheres than at 1 atmosphere [1]. The forward reaction turns 3 molecules into 1, and increased pressure favours the reaction producing fewer molecules [1]. (c) At 100 °C the reaction would be too slow [1]. Examiner insight: Quote figures from the table, such as 52% falling to 30% at 1 atmosphere.

5. Any six points, one mark each:

• Nitrogen comes from the air [1].
• Hydrogen comes from natural gas (or other sources) [1].
• The purified gases pass over an iron catalyst [1].
• at about 450 °C and about 200 atmospheres [1].
• Only some of the gases react because the reaction is reversible; on cooling, the ammonia turns to liquid and is taken off [1].
• The nitrogen and hydrogen that did not react go back into the reactor (recycled) [1]. Examiner insight: Follow the order raw materials → reactor → cooling → recycling, and give each condition with its unit.

6. (a) The forward reaction is exothermic, so a lower temperature would give a higher yield of ammonia [1]. However, the reaction would then be too slow to be worthwhile [1]. About 450 °C balances a reasonable yield against a fast enough rate (a compromise) [1]. (b) 4 molecules of gas become 2, so a high pressure increases the yield (and the rate) [1]. A much higher pressure would need stronger, more expensive equipment and more energy to compress the gases [1], and would be more dangerous [1]. Examiner insight: A “compromise” answer is strongest when both sides are stated: what is gained (yield or rate) and what is lost (rate, cost or safety).

7. (a) The world population is rising, so more food is needed from limited farmland [1]. Crops need nitrogen to grow (to make proteins), and harvesting removes nutrients from the soil [1]. Fertilizers replace these nutrients and so maximise food yields [1]. (b) Mᵣ: H₂SO₄ = 98, NH₃ = 17, (NH₄)₂SO₄ = 132. From the equation, 98 t of acid reacts with 2 × 17 = 34 t of ammonia [1]. 49 t is half of 98 t, so ammonia needed = 17 tonnes [1]. Ammonium sulfate formed = 132 ÷ 2 = 66 tonnes [1]. Examiner insight: Write the 2 : 1 ratio down; an answer of 8.5 tonnes of ammonia shows it was missed.

8. (a) Stage 1: S(s) + O₂(g) → SO₂(g) [1]. Stage 3: H₂O(l) + SO₃(g) → H₂SO₄(aq) [1]. (b) Vanadium(V) oxide, V₂O₅ [1]; around 450 °C [1]; atmospheric pressure [1]. Examiner insight: Give the full name with oxidation state or the formula V₂O₅; “vanadium” alone names a different substance and may not be accepted.

9. Any six points, one mark each:

• The V₂O₅ catalyst increases the rate of reaction [1].
• It does not change the yield, but equilibrium is reached faster [1].
• The forward reaction is exothermic, so a lower temperature would give a higher yield of SO₃ [1].
• But a lower temperature would make the reaction too slow, so about 450 °C is a compromise between rate and yield [1].
• 3 molecules become 2, so higher pressure would favour SO₃ [1].
• The yield is already high at atmospheric pressure, so the extra cost of high-pressure equipment and energy is not worth the small gain [1]. Examiner insight: Give each condition its own reason; merging temperature and pressure into one sentence blurs which reason is which.

10. (a) 15% of 560 kg = 84 kg of nitrogen [1]. (b) 84 ÷ 28 = 3 kmol N₂; this gives 2 × 3 = 6 kmol NH₃ [1]. Mass = 6 × 17 = 102 kg [1]. (c) 3 × 3 = 9 kmol H₂; mass = 9 × 2 = 18 kg [1]. (d) The remaining 476 kg of nitrogen is recycled back into the reactor (with unreacted hydrogen) [1]. So raw materials are not wasted and far more ammonia is made overall than one pass gives [1]. Examiner insight: Check that mass is conserved (84 + 18 = 102); if your figures do not balance, look again at the mole ratio before moving on.

11. (a) S + O₂ → SO₂: 32 t of sulfur gives 64 t of sulfur dioxide [1]. So 4.8 t gives 4.8 × 64 ÷ 32 = 9.6 tonnes [1]. (b) 2SO₂ + O₂: 2 × 64 = 128 t of SO₂ needs 32 t of O₂ [1]. So 9.6 t needs 9.6 × 32 ÷ 128 = 2.4 tonnes [1]. (c) 1 S gives 1 H₂SO₄, so 32 t of sulfur gives 98 t of acid [1]. 4.8 × 98 ÷ 32 = 14.7 tonnes [1]. (d) Stage 2 is reversible, so not all the sulfur dioxide is converted to sulfur trioxide (or some product is lost during separation) [1]. Examiner insight: Use 32 for O₂, not 16; using the atomic mass of oxygen halves the answer to part (b).

Where marks are usually lost

• Defining equilibrium without “closed system” or without “same rate in both directions”.
• Saying the amounts are equal at equilibrium.
• Predicting a temperature effect before working out which direction is exothermic.
• Counting atoms instead of molecules, or ignoring the numbers in front of formulae.
• Claiming a catalyst increases yield.
• Giving one side of a compromise only.
• Swapping the catalysts or pressures of the Haber and contact processes.
• Forgetting the 2 : 1 ratios in the ammonium sulfate and Stage 2 calculations.

Next steps

• Go over the key points again in the revision notes for this unit.
• Return to the full teaching in the study guide for this unit.
• Practise more reacting masses with the quantitative chemistry practice set.
• See the whole course on the OxfordAQA IGCSE Chemistry hub and tick off topics on the printable checklist.
• Try one of the free 10-minute diagnostics.
• Book a free trial class.

Official syllabus

OxfordAQA International GCSE Chemistry (9202) specification, Version 6.2, for exams May/June 2018 onwards, published by OxfordAQA. This practice set covers section 3.8.2 Factors affecting equilibrium and section 3.8.3 Production of ammonia and sulfuric acid.