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school. “Paper 42” means Paper 4, variant 2.
Practice questions
Original questions written by HomiLabs in the style of the exam, with answers and mark schemes. They are not
Cambridge past paper questions and have not been reviewed by Cambridge. Try each one before opening the answer.
Q1Nitrogen oxides, pollution & catalytic converters · Core · Paper 1 style · 1 mark
(a)Which pollutant is removed from car exhaust gases by reduction in a catalytic converter?[1]
Show answer
(a) C nitrogen monoxide
NO is reduced to N₂ (e.g. 2NO + 2CO → N₂ + 2CO₂). CO and hydrocarbons (A, B) are oxidised, not reduced; CO₂ (D) is not removed at all.
Q2Nitrogen oxides, pollution & catalytic converters · Core · Paper 2 style · 4 marks
Traffic on the main roads of Lahore releases oxides of nitrogen.
(a)Explain why nitrogen monoxide, NO, is formed inside a car engine although nitrogen and oxygen do not react in the air at room temperature.[2]
(b)Write an equation for the reaction of NO in the air to form NO₂.[1]
(c)State one natural source of oxides of nitrogen.[1]
Show answer
(a) The very high temperature in the engine (spark/combustion) provides enough energy to break the strong N≡N triple bond (overcome the high activation energy): N₂ + O₂ → 2NO.
B1 high temperature in engine. B1 enough energy to break N≡N / overcome high Ea (equation accepted).
(b) 2NO + O₂ → 2NO₂
B1
(c) Lightning (or volcanic activity / bacterial action in soil).
B1
Q3Nitrogen oxides, pollution & catalytic converters · Core · Paper 2 style · 5 marks
(a)Write the equation for the reaction between carbon monoxide and nitrogen monoxide in a catalytic converter.[1]
(b)Unburned octane, C₈H₁₈, can also be removed by reaction with NO, forming CO₂, H₂O and N₂. Write a balanced equation.[2]
(c)Name one metal used as the catalyst.[1]
(d)Explain why the catalyst is spread as a thin layer on a ceramic honeycomb.[1]
(d) Gives a very large surface area for the gases to be adsorbed, using little expensive metal.
B1
Q4Nitrogen oxides, pollution & catalytic converters · Core · Paper 2 style · 4 marks
An air-quality station beside a busy road in Lahore recorded these concentrations (in µg m⁻³):time | NO | NO₂06:00 | 20 | 1509:00 | 110 | 4012:00 | 40 | 85
(a)Suggest why the NO concentration rises sharply between 06:00 and 09:00.[1]
(b)Explain why NO₂ reaches its highest concentration later in the day than NO does. Include an equation.[2]
(c)Suggest why the NO concentration falls between 09:00 and 12:00.[1]
Show answer
(a) Morning rush-hour traffic: engines produce NO from N₂ and O₂.
B1
(b) NO is gradually oxidised by oxygen in the air: 2NO + O₂ → 2NO₂.
B1 equation. B1 NO converted to NO₂ over time.
(c) Less traffic after rush hour, and NO is being converted into NO₂ (and dispersed).
B1
Q5Nitrogen oxides, pollution & catalytic converters · Standard · Paper 1 style · 1 mark
(a)In sunlit air, nitrogen monoxide from vehicle exhausts is converted into nitrogen dioxide. Which row correctly gives the change in oxidation number of nitrogen and the colour of nitrogen dioxide?[1]
Show answer
(a) A +2 → +4; brown
In NO nitrogen is +2; in NO₂ it is +4 (2NO + O₂ → 2NO₂, an oxidation). NO₂ is a brown gas. B has the correct oxidation numbers but NO is the colourless gas, not NO₂; C reverses the change.
Q6Nitrogen oxides, pollution & catalytic converters · Standard · Paper 1 style · 1 mark
(a)In a catalytic converter: 2NO + 2CO → N₂ + 2CO₂What are the changes in oxidation number of nitrogen and of carbon?[1]
Show answer
(a) A N: +2 → 0; C: +2 → +4
In NO nitrogen is +2 and in N₂ it is 0 (reduced); carbon in CO is +2 and in CO₂ it is +4 (oxidised). B uses the value for NO₂; C reverses the carbon change.
Q7Nitrogen oxides, pollution & catalytic converters · Standard · Paper 2 style · 5 marks
(a)Write two equations to show how NO₂ catalyses the oxidation of SO₂ in the atmosphere.[2]
(b)Explain how your equations show that NO₂ acts as a catalyst.[1]
(c)Give the oxidation number of nitrogen in NO₂ and in NO.[1]
(d)Explain why this reaction makes the problem of acid rain worse.[1]
Show answer
(a) SO₂ + NO₂ → SO₃ + NO; NO + ½O₂ → NO₂
B1 each.
(b) NO₂ is used in the first step but re-formed in the second, so it is not used up overall (overall SO₂ + ½O₂ → SO₃).
B1
(c) +4 in NO₂; +2 in NO
B1
(d) SO₃ formed dissolves in rainwater to give sulfuric acid, H₂SO₄.
B1
Q8Nitrogen oxides, pollution & catalytic converters · Standard · Paper 2 style · 5 marks
A rickshaw engine without a catalytic converter releases 1.50 m³ of exhaust gas (measured at room conditions) for each kilometre travelled. The exhaust contains 0.12 % of NO by volume.
(a)Calculate the volume of NO, in dm³, released per kilometre.[1]
(b)Calculate the mass of NO released per kilometre. (molar gas volume 24.0 dm³ at room conditions; Mr NO = 30.0)[2]
(c)If a catalytic converter were fitted, state the volume of CO needed to remove all of this NO and the volume of N₂ formed, both at room conditions.[2]
Show answer
(a) 1.80 dm³
1.50 m³ = 1500 dm³; 0.12 % of 1500 = 1.80 dm³
(b) 2.25 g
M1 mol NO = 1.80 ÷ 24.0 = 0.0750. A1 × 30.0 = 2.25 g
(c) CO 1.80 dm³; N₂ 0.90 dm³
2CO + 2NO → 2CO₂ + N₂: CO : NO = 1 : 1 and N₂ : NO = 1 : 2.
Q9Nitrogen oxides, pollution & catalytic converters · Standard · Paper 2 style · 5 marks
Peroxyacetyl nitrate (PAN) is a component of photochemical smog. It causes eye irritation and damages plants.
(a)Name the two types of pollutant from vehicle exhausts that react together in sunlight to form PAN.[2]
(b)PAN has the molecular formula C₂H₃NO₅. Calculate its relative molecular mass.[1]
(c)Suggest why levels of photochemical smog in a large city such as Karachi are usually highest on sunny afternoons rather than early in the morning.[2]
Show answer
(a) Oxides of nitrogen (NO/NO₂) and unburned hydrocarbons
B1 each.
(b) 121.0
2(12.0) + 3(1.0) + 14.0 + 5(16.0) = 121.0
(c) Morning traffic first releases NOx and hydrocarbons; the reactions that form PAN need sunlight (UV) and time, so secondary pollutants build up later in the day.
B1 sunlight needed for the reactions. B1 time for primary pollutants from rush-hour traffic to react / build up.
Q10Nitrogen oxides, pollution & catalytic converters · Standard · Paper 2 style · 4 marks
Exhaust gas from a car engine was analysed before and after its catalytic converter (percentages by volume):gas | before | afterCO | 1.20 | 0.10NO | 0.080 | 0.012
(a)Calculate the percentage of CO and of NO removed by the converter. Give each answer to 3 significant figures.[2]
(b)Suggest why the converter removes much less pollution in the first few minutes after a cold start.[1]
(c)Explain why the percentage of CO₂ in the exhaust is higher after the converter.[1]
(b) The catalyst is cold, so the reactions are slow until it reaches its working temperature.
B1
(c) CO and hydrocarbons are oxidised to CO₂.
B1
Q11Nitrogen oxides, pollution & catalytic converters · Challenge · Paper 1 style · 1 mark
(a)Which statement about photochemical smog is correct?[1]
Show answer
(a) A Peroxyacetyl nitrate (PAN) forms when oxides of nitrogen react with unburned hydrocarbons.
PAN is a secondary pollutant made in the air from NO/NO₂ and hydrocarbons, so D is wrong. The converter removes NO (B wrong); the reactions need sunlight (C wrong).
Q12Nitrogen oxides, pollution & catalytic converters · Challenge · Paper 2 style · 6 marks
(a)Use Le Chatelier's principle to explain why more NO is present at the high temperatures inside an engine.[2]
(b)Explain why changing the pressure has no effect on the position of this equilibrium.[1]
(c)Write the expression for Kc and calculate the equilibrium concentration of NO at 2000 K when [N₂] = 4.8 × 10⁻³ and [O₂] = 1.2 × 10⁻³ mol dm⁻³. Give your answer to 2 significant figures.[3]
Show answer
(a) The forward reaction is endothermic, so increasing the temperature shifts the position of equilibrium to the right, increasing the yield of NO.
M1 forward reaction endothermic. A1 equilibrium moves right at high T.
(b) There are equal numbers of moles of gas on each side (2 and 2).
Q13Nitrogen oxides, pollution & catalytic converters · Challenge · Paper 2 style · 6 marks
(a)Describe, in terms of adsorption, how the metal catalyst in a catalytic converter speeds up the reaction between CO and NO.[3]
(b)Some 'lean-burn' engines run with a large excess of air. Suggest why a normal catalytic converter removes NO less effectively from these engines.[2]
(c)Explain why leaded petrol must not be used in cars with catalytic converters.[1]
Show answer
(a) Gas molecules are adsorbed onto the metal surface (weak bonds form to active sites); bonds within the molecules are weakened and the molecules are held close together in a favourable orientation, lowering Ea; the products (CO₂, N₂) are desorbed, freeing the sites.
B1 adsorption onto surface/active sites. B1 bonds weakened / molecules held close, so lower activation energy. B1 desorption of products.
(b) The excess O₂ oxidises CO and hydrocarbons on the catalyst, so little CO/hydrocarbon is left to reduce the NO to N₂.
B1 O₂ reacts with CO/hydrocarbons. B1 so not enough reducing agent for NO.
(c) Lead is adsorbed strongly onto the active sites and poisons the catalyst.
B1
Q14Ammonia & ammonium ion · Core · Paper 1 style · 1 mark
(a)What are the shape and bond angle of the ammonium ion, NH₄⁺?[1]
Show answer
(a) B tetrahedral, 109.5°
NH₄⁺ has four bonding pairs and no lone pair, so it is tetrahedral. A describes NH₃, which has one lone pair.
Q15Ammonia & ammonium ion · Core · Paper 2 style · 4 marks
(a)Describe the bonding in the ammonium ion, NH₄⁺, and explain how it forms from NH₃ and H⁺.[2]
(b)Give the H–N–H bond angle in NH₃ and in NH₄⁺, and explain the difference.[2]
Show answer
(a) Four covalent N–H bonds; the fourth is a dative (coordinate) bond formed when the lone pair on N is donated to H⁺. All four bonds are then identical.
B1 dative/coordinate bond. B1 lone pair on N donated to H⁺.
(b) NH₃ 107°, NH₄⁺ 109.5°. NH₃ has a lone pair, which repels more strongly than bonding pairs, pushing the N–H bonds closer; NH₄⁺ has four bonding pairs only.
B1 both angles. B1 lone pair–bond pair repulsion greater than bond pair–bond pair.
Q16Ammonia & ammonium ion · Core · Paper 2 style · 4 marks
A farmer near Multan buys a bag of white fertiliser that is labelled as an ammonium salt.
(a)Describe a test to show that a fertiliser sample contains ammonium ions. Give the reagent, condition and positive result.[3]
(b)Write the ionic equation for the reaction in the test.[1]
Show answer
(a) Add aqueous sodium hydroxide and warm; a gas is given off that turns damp red litmus paper blue (ammonia).
B1 NaOH(aq). B1 warm. B1 damp red litmus turns blue.
(b) NH₄⁺ + OH⁻ → NH₃ + H₂O
B1
Q17Ammonia & ammonium ion · Standard · Paper 1 style · 1 mark
(a)Which mixture gives off ammonia when warmed?[1]
Show answer
(a) A ammonium chloride and aqueous sodium hydroxide
OH⁻ removes H⁺ from NH₄⁺: NH₄⁺ + OH⁻ → NH₃ + H₂O. In D the only reaction is precipitation of BaSO₄; no base is present to release NH₃.
Q18Ammonia & ammonium ion · Standard · Paper 2 style · 6 marks
2.00 g of a fertiliser containing ammonium sulfate is boiled with 50.0 cm³ of 1.00 mol dm⁻³ NaOH (an excess) until no more ammonia is given off. The remaining NaOH needs 41.20 cm³ of 0.500 mol dm⁻³ HCl for neutralisation.
(a)Write the equation for the reaction of ammonium sulfate with sodium hydroxide.[1]
(b)Calculate the amount, in moles, of NaOH that reacted with the ammonium ions.[2]
(c)Calculate the percentage by mass of ammonium sulfate in the fertiliser. (Mr (NH₄)₂SO₄ = 132.1)[2]
(d)Calculate the percentage by mass of nitrogen in the fertiliser.[1]
Show answer
(a) (NH₄)₂SO₄ + 2NaOH → Na₂SO₄ + 2NH₃ + 2H₂O
B1
(b) 0.0294 mol
M1 NaOH added = 0.0500 mol; left over = 0.04120 × 0.500 = 0.0206. A1 0.0500 − 0.0206 = 0.0294
(b) It is only partially protonated (the equilibrium lies well to the left), so few OH⁻ ions form.
B1
(c) 2NH₃ + H₂SO₄ → (NH₄)₂SO₄
B1
Q20Ammonia & ammonium ion · Standard · Paper 3 style · 6 marks
A student heats a mixture of 1.07 g of ammonium chloride with excess calcium hydroxide and collects the ammonia in a gas syringe. 432 cm³ of ammonia is collected at room conditions.
(a)Write the equation for the reaction.[1]
(b)Calculate the maximum volume of ammonia, in cm³ at room conditions. (Mr NH₄Cl = 53.5; molar gas volume 24.0 dm³)[2]
(c)Calculate the percentage yield.[1]
(d)Explain why the gas is collected in a dry syringe rather than over water, and suggest one reason for the yield being below 100 %.[2]
(d) Ammonia is very soluble in water; yield low because the reaction was incomplete / some NH₃ escaped before the syringe was connected / NH₃ dissolved in condensed water in the tube.
B1 very soluble in water. B1 any valid reason.
Q21Ammonia & ammonium ion · Challenge · Paper 1 style · 1 mark
(a)In which reaction does the nitrogen-containing reactant act as a Brønsted–Lowry acid?[1]
Show answer
(a) B NH₄⁺ + OH⁻ → NH₃ + H₂O
NH₄⁺ donates a proton to OH⁻. In A and C ammonia accepts a proton (base); D is not an acid–base reaction.
Q22Ammonia & ammonium ion · Challenge · Paper 2 style · 5 marks
Ammonia is made at a fertiliser plant near Daharki by the Haber process:N₂(g) + 3H₂(g) ⇌ 2NH₃(g) ΔH = −92 kJ mol⁻¹In one equilibrium mixture at a total pressure of 20.0 MPa the mole fractions are N₂ 0.20, H₂ 0.60 and NH₃ 0.20.
(a)Write the expression for Kp and calculate its value, with units, to 3 significant figures.[3]
(b)Explain why the plant runs at about 450 °C even though a lower temperature gives a higher equilibrium yield.[2]
(b) The forward reaction is exothermic so a low temperature favours yield, but the rate would be too slow; 450 °C is a compromise giving acceptable yield at a fast enough rate (with an iron catalyst).
B1 lower T gives higher yield (exothermic). B1 rate too slow at low T; compromise.
Q23Ammonia & ammonium ion · Challenge · Paper 2 style · 5 marks
(a)Calculate the percentage by mass of nitrogen in ammonium nitrate, NH₄NO₃ (Mr 80.0), ammonium sulfate, (NH₄)₂SO₄ (Mr 132.1), and urea, CO(NH₂)₂ (Mr 60.0). Give each to 3 significant figures.[3]
(b)A farmer spreads slaked lime, Ca(OH)₂, on acidic soil at the same time as ammonium sulfate. Explain, with an equation, why some of the nitrogen is wasted.[2]
Show answer
(a) NH₄NO₃ 35.0 %; (NH₄)₂SO₄ 21.2 %; urea 46.7 %
B1 each: 28.0 ÷ Mr × 100
(b) (NH₄)₂SO₄ + Ca(OH)₂ → CaSO₄ + 2NH₃ + 2H₂O; the base displaces ammonia gas, which escapes into the air.
B1 equation. B1 NH₃ gas lost to the air.
Q24Sulfur dioxide & acid rain · Core · Paper 2 style · 5 marks
The Thar coalfield in Sindh supplies power stations that burn lignite.
(a)Explain how sulfur dioxide is released when coal is burned in a power station.[1]
(b)Write two equations to show how SO₂ is converted into sulfuric acid in the atmosphere.[2]
(c)State two harmful effects of acid rain.[2]
Show answer
(a) Coal contains sulfur compounds, which burn to form SO₂ (S + O₂ → SO₂).
B1
(b) SO₂ + ½O₂ → SO₃; SO₃ + H₂O → H₂SO₄
B1 each (accept SO₂ + H₂O → H₂SO₃ then oxidation).
(c) Any two: kills fish/aquatic life in lakes; damages trees/leaches nutrients from soil; corrodes limestone/marble buildings and metal structures; releases toxic Al³⁺ into water.
1 mark each for any two.
Q25Sulfur dioxide & acid rain · Standard · Paper 2 style · 5 marks
A coal sample contains 1.8 % sulfur by mass. Assume all the sulfur is converted to SO₂ on burning. (Ar S = 32.1)
(a)Calculate the amount, in moles, of sulfur in 1.00 tonne (1.00 × 10⁶ g) of coal.[1]
(b)Calculate the volume of SO₂, in m³ at room conditions, formed when 1.00 tonne of coal burns. (molar gas volume 24.0 dm³)[1]
(c)The SO₂ is removed by powdered calcium carbonate: CaCO₃ + SO₂ → CaSO₃ + CO₂. Calculate the mass of CaCO₃, in kg, needed per tonne of coal. (Mr CaCO₃ = 100.1)[2]
(d)Suggest one disadvantage of this method of removing SO₂.[1]
Show answer
(a) 561 mol
1.8 % of 10⁶ g = 18 000 g; ÷ 32.1 = 561 mol
(b) 13.5 m³
560.7 × 24.0 dm³ = 13458 dm³ = 13.5 m³
(c) 56.1 kg
M1 mol CaCO₃ = mol SO₂ = 560.7. A1 × 100.1 = 56131 g = 56.1 kg
(d) It releases CO₂ (a greenhouse gas) / limestone must be quarried / large amounts of solid waste.
B1
Q26Sulfur dioxide & acid rain · Standard · Paper 2 style · 4 marks
(a)Write the equation for the reaction of sulfuric acid in rain with marble, CaCO₃.[1]
(b)During one year, 5000 dm³ of rain containing 2.0 × 10⁻⁵ mol dm⁻³ of H₂SO₄ falls on a marble monument. Calculate the maximum mass of marble that can be dissolved. (Mr CaCO₃ = 100.1)[2]
Q27Sulfur dioxide & acid rain · Challenge · Paper 2 style · 6 marks
To monitor sulfur dioxide near a brick kiln, 50.0 m³ of air is bubbled through excess hydrogen peroxide solution, which converts all the SO₂ to sulfuric acid. The solution needs 18.60 cm³ of 0.0100 mol dm⁻³ NaOH for neutralisation. (Mr SO₂ = 64.1)
(a)Write the equation for the reaction of SO₂ with H₂O₂ and give the change in oxidation number of sulfur.[2]
(b)Calculate the mass of SO₂ in the air sample, in µg, to 3 significant figures.[3]
(c)Calculate the SO₂ concentration in µg m⁻³ and state whether it is above a guideline limit of 40 µg m⁻³ (24-hour average).[1]
Q28Nitrogen: unreactivity · Core · Paper 2 style · 4 marks
(a)Explain why nitrogen gas is very unreactive.[2]
(b)Give one use of nitrogen that depends on its lack of reactivity.[1]
(c)The bond energy of O=O is 496 kJ mol⁻¹ and that of N≡N is 944 kJ mol⁻¹. Oxygen is far more reactive than nitrogen. Use the bond energies to explain this difference.[1]
Show answer
(a) The N≡N triple bond is very strong (high bond energy), and the molecule is non-polar, so it does not attract electrophiles or nucleophiles.
B1 strong triple bond / high bond energy. B1 non-polar molecule.
(b) Filling packets of food (e.g. crisps) to prevent oxidation / an inert atmosphere for welding or storing chemicals.
B1
(c) Much less energy is needed to break O=O than N≡N, so reactions of O₂ have lower activation energies.
B1
Q29Nitrogen: unreactivity · Standard · Paper 2 style · 5 marks
(b) The activation energy is very high because the N≡N bond must be broken first; an iron catalyst (and a raised temperature) is used to give a reasonable rate.
B1 high activation energy (N≡N). B1 iron catalyst / high temperature.
Q30Nitric acid & nitrates · Standard · Paper 2 style · 5 marks
Ammonium nitrate fertiliser is made in Pakistan from ammonia, part of which is oxidised to nitric acid.
(a)Write the equation for the formation of ammonium nitrate from ammonia and nitric acid.[1]
(b)Give the oxidation number of each nitrogen atom in ammonium nitrate.[1]
(c)In the factory, half of the ammonia is first converted into nitric acid, which then reacts with the rest of the ammonia. Calculate the maximum mass of ammonium nitrate, in tonnes, made from 1.00 tonne of ammonia. (Mr NH₃ = 17.0; Mr NH₄NO₃ = 80.0)[3]
