2022 May/June · Paper 42 · Q4 (b(ii))4 marksTests for ions, gases & water
2022 May/June · Paper 42 · Q5 (a, b, c(i))3 marksApparatus, measurement & gas collection · Practical planning & evaluation · Results tables, graphs & data analysis
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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.
Q1Tests for ions, gases & water · Core · Paper 1 style · 1 mark
(a)[1]
Show answer
(a) A Al³⁺
Al³⁺: white ppt soluble in excess NaOH but insoluble in excess NH₃. Zn²⁺ (B) is wrong because its white ppt also dissolves in excess aqueous ammonia; Ca²⁺ ppt does not dissolve in excess NaOH; Cu²⁺ gives a light blue ppt.
Q2Tests for ions, gases & water · Core · Paper 1 style · 1 mark
(a)[1]
Show answer
(a) B sulfur dioxide
Sulfur dioxide is a reducing agent and decolourises acidified manganate(VII). Carbon dioxide (A) turns limewater milky but has no effect on manganate(VII).
Q3Tests for ions, gases & water · Core · Paper 2 style · 5 marks
(a)A family in Karachi collects water from a well. Describe a chemical test to show that a sample contains water. Give the result.[2]
(b)Describe a physical test to show whether the water is pure.[2]
(c)A sample of the well water is acidified with dilute nitric acid and aqueous silver nitrate is added. A white precipitate forms. Name the anion shown to be present.[1]
Show answer
(a) Add anhydrous cobalt(II) chloride; it turns from blue to pink (or anhydrous copper(II) sulfate turns from white to blue).
B1 correct reagent (anhydrous CoCl₂ or anhydrous CuSO₄); B1 correct colour change from → to
(b) Measure its boiling point (or melting point); pure water boils at exactly 100 °C (melts at exactly 0 °C).
B1 measure boiling point / melting point; B1 100 °C / 0 °C (at standard pressure)
(c) chloride ion, Cl⁻
B1
Q4Tests for ions, gases & water · Core · Paper 2 style · 5 marks
(a)Aqueous ammonia is added a drop at a time to copper(II) sulfate solution until it is in excess. Describe what is seen.[2]
(b)Describe the test for carbon dioxide and give the result.[2]
(c)State the colour of the precipitate formed when aqueous sodium hydroxide is added to a solution containing iron(III) ions.[1]
Show answer
(a) A light blue precipitate forms, which dissolves in excess ammonia to give a dark blue solution.
B1 light blue precipitate; B1 dissolves in excess to dark blue solution
(b) Bubble the gas through limewater; it turns milky (cloudy white).
B1 limewater; B1 milky / cloudy
(c) red-brown
B1
Q5Tests for ions, gases & water · Standard · Paper 1 style · 1 mark
(a)Which reagents are used to show that a solution contains sulfate ions?[1]
Show answer
(a) B dilute nitric acid, then aqueous barium nitrate
Sulfate gives a white precipitate (barium sulfate) with aqueous barium nitrate in acid solution. A tests for halide ions; C tests for nitrate; D tests for sulfite.
Q6Tests for ions, gases & water · Standard · Paper 4 style · 5 marks
A student in a school laboratory in Multan is given a pale green crystalline salt X. X is dissolved in distilled water and the solution is divided into three portions.Test 1: aqueous sodium hydroxide is added until in excess. Observation: green precipitate, insoluble in excess; on standing the surface of the precipitate turns brown.Test 2: aqueous ammonia is added until in excess. Observation: green precipitate, insoluble in excess.Test 3: the solution is acidified with dilute nitric acid and aqueous barium nitrate is added. Observation: white precipitate.
(a)Identify the cation in X. Explain your answer using test 1.[2]
(b)Suggest why the precipitate in test 1 turns brown at the surface on standing.[1]
(c)Identify the anion in X.[1]
(d)Name salt X.[1]
Show answer
(a) Iron(II), Fe²⁺: green precipitate that does not dissolve in excess sodium hydroxide.
B1 Fe²⁺; B1 green ppt insoluble in excess (must be green, not just 'precipitate')
(b) Fe²⁺ (iron(II) hydroxide) is oxidised by oxygen in the air to iron(III) hydroxide, which is red-brown.
B1 oxidation by air / oxygen to iron(III)
(c) sulfate ion, SO₄²⁻
B1
(d) iron(II) sulfate
B1 (allow FeSO₄)
Q7Tests for ions, gases & water · Standard · Paper 3 style · 5 marks
A laboratory technician in Peshawar finds three jars of solid salts, P, Q and R, with damaged labels. Each is one of: calcium carbonate, potassium bromide, copper(II) sulfate.P: lilac flame colour; its solution acidified with dilute nitric acid gives a cream precipitate with aqueous silver nitrate.Q: orange-red flame colour; with dilute hydrochloric acid it effervesces and the gas turns limewater milky.R: blue-green flame colour; its solution gives a light blue precipitate with aqueous sodium hydroxide, and a white precipitate with dilute nitric acid and aqueous barium nitrate.
(a)Name salt P.[1]
(b)Name salt Q. Name the gas given off with the acid.[2]
(c)Name salt R.[1]
(d)State the flame colour expected for a sodium salt.[1]
Show answer
(a) potassium bromide
B1 (K⁺ from lilac flame, Br⁻ from cream ppt)
(b) calcium carbonate; carbon dioxide
B1 calcium carbonate; B1 carbon dioxide
(c) copper(II) sulfate
B1 (Cu²⁺ from blue-green flame and light blue ppt; SO₄²⁻ from white ppt with barium nitrate)
(d) yellow
B1
Q8Tests for ions, gases & water · Standard · Paper 2 style · 8 marks
A student carries out several reactions that produce gases.
(a)Describe the test for hydrogen and give the result.[2]
(b)Describe the test for oxygen and give the result.[2]
(c)A gas from a reaction bleaches damp litmus paper. Name the gas.[1]
(d)Ammonium chloride is warmed with aqueous sodium hydroxide. Name the gas formed and give a test for it with the result.[3]
Show answer
(a) Hold a lighted splint at the mouth of the tube; it gives a squeaky pop.
B1 lighted (burning) splint; B1 pops
(b) Insert a glowing splint; it relights.
B1 glowing splint; B1 relights
(c) chlorine
B1
(d) Ammonia; damp red litmus paper turns blue.
B1 ammonia; B1 damp red litmus; B1 turns blue
Q9Tests for ions, gases & water · Challenge · Paper 4 style · 6 marks
Solid Y gives a yellow flame colour. A solution of Y is tested.Step 1: aqueous sodium hydroxide is added. No precipitate forms.Step 2: the mixture is warmed. No gas is given off.Step 3: a small piece of aluminium foil is added and the mixture is warmed again. A gas is given off that turns damp red litmus paper blue.
(a)Identify the cation in Y.[1]
(b)Explain why step 2 (warming with sodium hydroxide before adding aluminium) is needed.[2]
(c)Name the gas given off in step 3 and identify the anion in Y.[2]
(d)Name compound Y.[1]
Show answer
(a) sodium ion, Na⁺
B1 (yellow flame; no precipitate with NaOH)
(b) To show that ammonium ions are absent, so that any ammonia formed in step 3 must come from nitrate ions being reduced.
B1 shows no ammonium ions present (NH₄⁺ would give ammonia with NaOH alone); B1 so ammonia in step 3 comes from nitrate
(c) ammonia; nitrate ion, NO₃⁻
B1 ammonia; B1 nitrate
(d) sodium nitrate
B1
Q10Tests for ions, gases & water · Challenge · Paper 2 style · 6 marks
(a)Two colourless solutions are known to contain Zn²⁺ ions and Al³⁺ ions, but they are unlabelled. Both give a white precipitate that dissolves in excess aqueous sodium hydroxide. Describe a test that distinguishes them, giving the result for each.[3]
(b)Describe how a solution of potassium iodide could be distinguished from a solution of potassium chloride using one reagent after acidifying. Give both results.[2]
(c)Explain why the solutions in (b) are acidified with dilute nitric acid and not with dilute hydrochloric acid.[1]
Show answer
(a) Add aqueous ammonia until in excess: Zn²⁺ gives a white precipitate that dissolves in excess (colourless solution); Al³⁺ gives a white precipitate that does not dissolve in excess.
(b) Add aqueous silver nitrate: iodide gives a yellow precipitate; chloride gives a white precipitate.
B1 silver nitrate with yellow ppt for iodide; B1 white ppt for chloride
(c) Hydrochloric acid contains chloride ions, which would give a white precipitate with silver nitrate whatever the sample.
B1 HCl adds Cl⁻ ions (false positive)
Q11Tests for ions, gases & water · Challenge · Paper 4 style · 6 marks
Two solutions, S and T, are tested.Solution S, test 1: aqueous sodium hydroxide added until in excess. Green precipitate, soluble in excess giving a green solution.Solution S, test 2: aqueous ammonia added until in excess. Green precipitate, insoluble in excess.Solution T, test 3: acidified aqueous potassium manganate(VII) added. The purple colour disappears.Solution T, test 4: acidified with dilute nitric acid, then aqueous barium nitrate added. No precipitate.
(a)Identify the cation in solution S. Explain how test 1 and test 2 together rule out iron(II) ions.[3]
(b)Identify the anion in solution T. Explain your answer.[2]
(c)Explain why test 4 shows that T does not contain sulfate ions.[1]
Show answer
(a) Chromium(III), Cr³⁺. The green precipitate dissolves in excess sodium hydroxide; the green iron(II) hydroxide precipitate does not dissolve in excess.
B1 Cr³⁺; B1 green ppt soluble in excess NaOH; B1 Fe²⁺ ppt insoluble in excess NaOH
(b) Sulfite, SO₃²⁻; it decolourises acidified potassium manganate(VII) (purple to colourless).
B1 sulfite; B1 purple → colourless with acidified KMnO₄
(c) Sulfate ions would give a white precipitate (barium sulfate) with acidified aqueous barium nitrate; no precipitate formed.
B1 sulfate would give white ppt of barium sulfate; none seen
Q12Practical planning & evaluation · Core · Paper 4 style · 4 marks
A student in Faisalabad mixes 25.0 cm³ of dilute hydrochloric acid and 25.0 cm³ of aqueous sodium hydroxide in a glass beaker. She measures the temperature before mixing and the highest temperature after mixing with a thermometer reading to the nearest 1 °C.
(a)Suggest why the temperature rise measured was smaller than expected.[1]
(b)Suggest two improvements to the apparatus or method.[2]
(c)Explain why the mixture should be stirred before the highest temperature is read.[1]
Show answer
(a) Heat (thermal energy) was lost to the surroundings through the glass beaker.
B1 heat loss to surroundings / glass is a good conductor
(b) Use a polystyrene cup (insulated) instead of a glass beaker; put a lid on the cup; (or use a thermometer reading to 0.1 °C / read at eye level).
1 mark each for any two of: polystyrene/insulated cup; lid; thermometer with finer scale or digital; read thermometer at eye level
(c) So that the reactants mix fully and the temperature is the same throughout.
B1 even temperature / complete mixing
Q13Practical planning & evaluation · Standard · Paper 3 style · 6 marks
A pharmacy in Islamabad sells three brands of antacid tablet, A, B and C. Each brand contains a base that neutralises stomach acid. The tablets are all the same mass.
(a)Plan an experiment to find which tablet neutralises the most hydrochloric acid. You are provided with dilute hydrochloric acid of known concentration, a burette, methyl orange indicator and common laboratory apparatus. In your plan include how you make the comparison fair and how you decide which tablet is best.[6]
Show answer
(a) Crush one tablet of A and dissolve/suspend it in a fixed volume of distilled water in a conical flask; add a few drops of methyl orange; titrate with the hydrochloric acid from the burette, swirling, until the colour changes from yellow to orange (permanent); record the volume of acid added; repeat with one tablet of B and of C using the same mass/number of tablets and same acid; repeat each until results are concordant; the tablet needing the largest volume of acid neutralises the most.
1 mark each for any six of: crush tablet (or same mass of each powder); add to fixed volume of water in conical flask; add methyl orange; add acid from burette until colour change yellow → orange/red; record volume of acid used; same acid concentration / same mass or one whole tablet each; repeat for concordant results / average; largest volume of acid = best
Q14Practical planning & evaluation · Standard · Paper 3 style · 6 marks
When dilute hydrochloric acid is added to sodium thiosulfate solution, a pale yellow precipitate of sulfur slowly forms and the mixture becomes cloudy.
(a)Plan an experiment to find how temperature affects the rate of this reaction. Include the variables you keep constant, how you change the temperature and what you measure.[6]
Show answer
(a) Draw a black cross on paper and place a conical flask on it; measure a fixed volume (e.g. 50 cm³) of sodium thiosulfate solution into the flask; warm it in a water bath to a chosen temperature and measure with a thermometer; add a fixed volume (e.g. 5 cm³) of hydrochloric acid of fixed concentration and start a stop-watch; time how long until the cross can no longer be seen; repeat at four or five other temperatures (e.g. 20–60 °C); keep volumes, concentrations, and the same cross / observer constant; shorter time = faster rate (rate ∝ 1/time).
1 mark each for any six of: cross under flask; fixed volume of thiosulfate; heat in water bath to set temperatures; measure temperature with thermometer; add fixed volume/concentration of acid and start timer; time until cross disappears; at least 5 different temperatures; same concentrations / volumes / cross kept constant; rate = 1/time (shorter time = faster)
Q15Practical planning & evaluation · Standard · Paper 3 style · 6 marks
(a)Plan an experiment to find the solubility of potassium nitrate at 40 °C, in grams per 100 g of water. Include how you make sure the solution is saturated and what you weigh.[6]
Show answer
(a) Heat distilled water in a beaker in a water bath at 40 °C (thermometer); add potassium nitrate a little at a time with stirring until some solid will no longer dissolve (saturated); filter or decant some of the saturated solution at 40 °C into a weighed evaporating basin; weigh basin + solution; evaporate the water (heat to dryness or to constant mass); weigh basin + solid; mass of water = mass of solution − mass of solid; solubility = mass of solid ÷ mass of water × 100.
1 mark each for any six of: water bath at 40 °C with thermometer; add solid until excess remains (saturated); remove undissolved solid (filter/decant) at 40 °C; weigh empty basin; weigh basin + solution; evaporate all the water / heat to constant mass; weigh basin + dry solid; calculation per 100 g water
Q16Practical planning & evaluation · Challenge · Paper 3 style · 6 marks
(a)Plan an experiment to find the percentage by mass of water in hydrated magnesium sulfate crystals. Include the measurements you would make, how you would make sure all the water is removed, and how you would calculate the result.[6]
Show answer
(a) Weigh an empty dry crucible; add crystals and weigh again; heat gently (with lid partly on) then more strongly; cool in a desiccator and weigh; repeat heating, cooling and weighing until two masses agree (constant mass); mass of water = mass before − mass after heating; % water = mass of water ÷ mass of hydrated crystals × 100.
1 mark each for any six of: weigh empty crucible; weigh crucible + crystals; heat (gently at first to avoid spitting); cool before weighing (desiccator); reweigh; heat to constant mass (repeat until mass unchanged); mass of water = loss in mass; % = loss ÷ mass of crystals × 100
Q17Practical planning & evaluation · Challenge · Paper 3 style · 6 marks
(a)Plan an experiment, using aqueous copper(II) sulfate and powders of zinc and magnesium, to find which metal is more reactive. Include the variables you would control and how you would reach a conclusion.[6]
Show answer
(a) Measure a fixed volume (e.g. 25 cm³) of copper(II) sulfate solution of fixed concentration into a polystyrene cup; record its starting temperature; add a fixed mass (e.g. 1.0 g, an excess) of zinc powder, stir, and record the highest temperature; repeat with fresh solution and the same mass of magnesium powder; keep volume, concentration, particle size and starting temperature the same; repeat each experiment; the metal giving the larger temperature rise is more reactive.
1 mark each for any six of: fixed volume of CuSO₄ in insulated cup; measure initial temperature; add fixed mass of metal powder; stir and record maximum temperature; repeat with other metal and fresh solution; same concentration / particle size / starting temperature; repeat for reliability; larger temperature rise = more reactive metal
Q18Titration technique · Core · Paper 1 style · 1 mark
(a)[1]
Show answer
(a) B 23.40 cm³
Burette readings are recorded to 2 decimal places (to the nearest 0.05 cm³). 23.4 cm³ (A) has only 1 decimal place.
Q19Titration technique · Core · Paper 4 style · 5 marks
25.0 cm³ of aqueous sodium carbonate is titrated with dilute hydrochloric acid using methyl orange indicator.Rough: initial 0.00, final 24.80Titration 1: initial 1.25, final 25.40Titration 2: initial 0.60, final 25.10Titration 3: initial 0.00, final 24.05(burette readings in cm³)
(a)Calculate the volume of acid added in each titration.[2]
(b)Which titrations are concordant (within 0.10 cm³ of each other)?[1]
(c)Calculate the average volume of acid to use in calculations.[1]
(d)State the colour of methyl orange at the end-point.[1]
B2 all four correct to 2 d.p.; B1 for three correct
(b) titrations 1 and 3
B1
(c) 24.10 cm³
(24.15 + 24.05) ÷ 2 = 24.10 cm³ B1
(d) orange
B1 (yellow in alkali, red in acid)
Q20Titration technique · Standard · Paper 2 style · 6 marks
A student in Quetta uses a pipette to put 25.0 cm³ of aqueous potassium hydroxide into a conical flask. She titrates it with dilute hydrochloric acid from a burette.
(a)State why the burette is rinsed with a little of the acid, not with water, before it is filled.[1]
(b)State why the conical flask is placed on a white tile.[1]
(c)Describe how the acid is added as the end-point is approached.[2]
(d)Thymolphthalein is used as the indicator, with the alkali in the flask. State the colour change at the end-point.[2]
Show answer
(a) Water left in the burette would dilute the acid and change its concentration.
B1 water would dilute the acid
(b) So that the colour change of the indicator is easier to see.
B1
(c) Drop by drop, swirling the flask after each drop, stopping when one drop causes a permanent colour change.
B1 dropwise; B1 swirl / stop at first permanent colour change
(d) blue to colourless
B1 blue (start); B1 colourless (end)
Q21Titration technique · Challenge · Paper 4 style · 8 marks
In a titration, aqueous sodium hydroxide from a pipette is in the conical flask and dilute sulfuric acid is added from the burette. For each mistake below, state whether the titre would be larger, smaller or unchanged, and explain why.
(a)Mistake 1: an air bubble in the burette tip escapes during the titration.[2]
(b)Mistake 2: the pipette is rinsed with distilled water just before measuring out the alkali.[2]
(c)Mistake 3: some distilled water is added to the conical flask to wash the sides during the titration.[2]
(d)Mistake 4: the student reads the burette from above the meniscus (looking down) at the end, but at eye level at the start.[2]
Show answer
(a) Titre recorded too large; the volume of the bubble is read as acid added.
B1 larger; B1 reason
(b) Titre too small; the alkali is diluted, so fewer moles of alkali are in the flask and less acid is needed.
B1 smaller; B1 alkali diluted / fewer moles
(c) No change; the number of moles of alkali in the flask is unchanged.
B1 no change; B1 moles unchanged
(d) Titre too small; viewed from above, the final reading is lower than the true value (burette scale numbers increase downwards), so final − initial is too small.
B1 smaller; B1 final reading too low (parallax error)
Q22Results tables, graphs & data analysis · Core · Paper 4 style · 6 marks
A student adds magnesium ribbon to excess dilute hydrochloric acid and collects the hydrogen in a gas syringe.Time / s: 0, 10, 20, 30, 40, 50, 60, 70Volume / cm³: 0, 28, 50, 66, 70, 84, 88, 88
(a)Identify the anomalous result and suggest the value it should be close to.[2]
(b)State the time by which the reaction had stopped and the total volume of gas collected.[2]
(c)Calculate the average rate of reaction in the first 30 s, in cm³/s.[2]
Show answer
(a) 40 s (70 cm³); about 77 cm³ (accept 75–79 cm³)
B1 40 s reading; B1 value in range 75–79
(b) 60 s; 88 cm³
B1 60 s; B1 88 cm³
(c) 2.2 cm³/s
66 ÷ 30 M1 = 2.2 cm³/s A1
Q23Results tables, graphs & data analysis · Standard · Paper 4 style · 5 marks
In a thermometric titration, aqueous sodium hydroxide is added in portions from a burette to 25.0 cm³ of dilute hydrochloric acid in a polystyrene cup. The temperature is recorded after each addition.Volume of NaOH / cm³: 0, 5.0, 10.0, 15.0, 25.0, 30.0, 35.0, 40.0Temperature / °C: 20.0, 23.0, 26.0, 29.0, 30.5, 29.0, 27.5, 26.0
(a)On a graph, two straight lines can be drawn through the points. Find the volume of sodium hydroxide at which the lines cross, and the highest temperature reached.[2]
(b)Explain why the temperature falls after this volume.[1]
(c)The hydrochloric acid has a concentration of 1.00 mol/dm³. Calculate the concentration of the sodium hydroxide in mol/dm³.[2]
Show answer
(a) 20.0 cm³; 32.0 °C
B1 rising line T = 20 + 0.6V, falling line T = 32 − 0.3(V − 20) → cross at 20.0 cm³; B1 32.0 °C
(b) Neutralisation is complete, so no more heat is released; the cooler alkali added (and heat loss) lowers the temperature.
Q24Results tables, graphs & data analysis · Standard · Paper 4 style · 6 marks
Calcium carbonate is added to dilute hydrochloric acid in a flask on a balance. A cotton-wool plug is placed in the neck.Time / min: 0, 1, 2, 3, 4, 5, 6Mass / g: 152.40, 151.62, 151.18, 150.96, 151.10, 150.80, 150.80
(a)Explain why the mass of the flask and contents decreases.[1]
(b)Identify the anomalous reading.[1]
(c)Calculate the total mass of carbon dioxide given off.[1]
(d)Explain why the mass stops changing after 5 minutes.[1]
(e)The acid was in excess. Calculate the mass of calcium carbonate that reacted. CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂ [Mr: CaCO₃, 100; CO₂, 44][2]
Show answer
(a) Carbon dioxide gas is produced and escapes from the flask.
B1
(b) 151.10 g at 4 min
B1 (mass cannot increase)
(c) 1.60 g
152.40 − 150.80 = 1.60 g B1
(d) The reaction has finished because one reactant (the calcium carbonate or the acid) has been used up.
B1
(e) 3.64 g
moles CO₂ = 1.60 ÷ 44 = 0.0364 M1; mass CaCO₃ = 0.0364 × 100 = 3.64 g A1
Q25Results tables, graphs & data analysis · Challenge · Paper 4 style · 6 marks
A student investigates how the concentration of sodium thiosulfate affects the time for a cross under the flask to disappear when acid is added. All other conditions are kept the same.Concentration / mol/dm³: 0.04, 0.08, 0.12, 0.16, 0.20Time / s: 125, 62, 42, 31, 25
(a)Calculate 1/time for the concentrations 0.04 and 0.20 mol/dm³. Give units.[2]
(b)Use 1/time to describe the relationship between concentration and rate.[2]
(c)Predict the time taken at 0.10 mol/dm³.[1]
(d)Suggest one reason why the times are not exact multiples of each other.[1]
Show answer
(a) 0.0080 s⁻¹ (/s) and 0.040 s⁻¹ (/s)
1 ÷ 125 = 0.0080 B1; 1 ÷ 25 = 0.040 B1 (units /s)
(b) Rate (1/time) is directly proportional to concentration: doubling the concentration doubles 1/time.
B1 rate increases with concentration; B1 directly proportional with evidence (e.g. 0.04 → 0.08 doubles 1/time)
(c) 50 s (accept 48–52 s)
c × t ≈ 5.0, so t = 5.0 ÷ 0.10 = 50 s B1
(d) Judging when the cross disappears is subjective / reaction time in starting and stopping the timer / small temperature changes.
B1 any sensible source of error
Q26Apparatus, measurement & gas collection · Core · Paper 1 style · 1 mark
(a)[1]
Show answer
(a) C a volumetric pipette
A volumetric pipette measures a fixed volume very accurately (25.0 cm³). A measuring cylinder (B) is less accurate.
Q27Apparatus, measurement & gas collection · Standard · Paper 2 style · 6 marks
(a)Hydrogen can be collected over water. State the property of hydrogen that allows this.[1]
(b)Explain why ammonia is not collected over water, and describe how it can be collected instead.[2]
(c)Chlorine is denser than air. Describe how a gas jar of chlorine can be collected.[1]
(d)Name the apparatus used to measure the volume of a gas produced in a rate experiment, and state the precision it is usually read to.[2]
Show answer
(a) It is insoluble (only very slightly soluble) in water.
B1
(b) Ammonia is very soluble in water. It is collected in an inverted gas jar/tube (upward delivery) because it is less dense than air.
B1 very soluble; B1 upward delivery / inverted container because less dense than air
(c) Downward delivery: the delivery tube goes to the bottom of an upright gas jar so chlorine pushes the air out of the top.
B1
(d) Gas syringe; to the nearest 1 cm³.
B1 gas syringe (allow inverted burette/measuring cylinder over water); B1 1 cm³
Q28Separation, purification & purity · Standard · Paper 1 style · 1 mark
(a)[1]
Show answer
(a) D fractional distillation
Two miscible liquids with fairly close boiling points are separated by fractional distillation using a fractionating column. Simple distillation (C) is used to separate a solvent from a dissolved solid.
Q29Separation, purification & purity · Challenge · Paper 2 style · 6 marks
(a)Rock salt from the Khewra mine is a mixture of sodium chloride and insoluble impurities such as clay and sand. Describe how a student could obtain pure dry sodium chloride crystals from rock salt.[4]
(b)A student tests the purity of a solid by finding its melting point. State how the melting point of an impure sample differs from that of the pure substance.[2]
Show answer
(a) Crush the rock salt and stir it in warm water so the sodium chloride dissolves; filter to remove the insoluble impurities (residue); heat the filtrate to evaporate some water until crystals start to form (saturated); leave to cool and crystallise; filter off the crystals, wash with a little cold water and dry between filter papers.
1 mark each for: dissolve in water (stir/warm); filter, impurities as residue; evaporate filtrate to point of crystallisation and cool; collect and dry crystals
(b) The impure sample melts at a lower temperature and over a range of temperatures (pure: sharp, fixed melting point).
B1 lower; B1 over a range (not sharp)
Q30Chromatography · Standard · Paper 4 style · 5 marks
A food inspector in Lahore tests the colouring from a sweet using paper chromatography. A spot of the colouring is placed on a pencil line 1.0 cm from the bottom of the paper. After the experiment, the solvent front is 8.0 cm above the starting line and two spots are seen, 2.4 cm and 5.6 cm above the starting line.Rf values of permitted dyes in the same solvent: dye E 0.30, dye F 0.55, dye G 0.70, dye H 0.85.
(a)Explain why the starting line is drawn in pencil and not in ink.[1]
(b)Calculate the Rf values of the two spots.[2]
(c)Use the data to identify which permitted dyes are in the sweet colouring.[1]
(d)Suggest how the spots of a colourless substance could be seen on a chromatogram.[1]
Show answer
(a) Pencil (graphite) does not dissolve in the solvent and move up the paper; ink would.
B1
(b) 0.30 and 0.70
Rf = distance moved by spot ÷ distance moved by solvent M1; 2.4 ÷ 8.0 = 0.30, 5.6 ÷ 8.0 = 0.70 A1
(c) E and G
B1
(d) Spray with a locating agent (then dry/heat).
B1
Similar questions
Questions that test the same subtopic in different years, grouped so you can work on one skill at a time.
References only: find each question in its paper (official link where available, otherwise through your school).
Tests for ions, gases & waterBanker81 questions · 9/9 sessions
