O Level Chemistry Practice Paper 1

TuitionGoWhere Practice Paper - Chemistry O-Level

TuitionGoWhere Secondary School (AI)

Subject: Chemistry
Level: O-Level
Paper: PRACTICE Paper 2
Duration: 1 hour 45 minutes
Total Marks: 80 marks

Name: _________________ Class: _________ Date: _________

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Instructions to Candidates

• Answer all questions in the spaces provided
• Show all working clearly for calculations
• The use of calculators is permitted
• Chemical formulae and equations should be written clearly
• Where numerical answers are required, give your answers to an appropriate number of significant figures

Information for Candidates

• The number of marks is given in brackets [ ] at the end of each question or part question
• You may find the Periodic Table useful
• At room temperature and pressure, 1 mole of any gas occupies 24 dm³

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Section A [40 marks]

1. A student investigates the properties of acids and bases.

(a) Complete the following statements about acids and bases: [2]

(i) Acids produce _____________ ions when dissolved in water.

(ii) Alkalis produce _____________ ions when dissolved in water.

(b) The student tests four solutions, P, Q, R and S, with Universal Indicator. The results are shown in the table below.

Solution	Colour with Universal Indicator	pH
P	Red	2
Q	Yellow	5
R	Green	7
S	Blue	9

(i) Which solution is neutral? [1]

(ii) Which solution is the strongest acid? Explain your answer. [2]

(iii) Arrange solutions P, Q and S in order of increasing H⁺ ion concentration. [1]

(c) The student adds magnesium ribbon to solutions P and Q separately.

(i) Write a word equation for the reaction between magnesium and an acid. [1]

(ii) State two observations the student would make when magnesium is added to solution P. [2]

(iii) Compare the rate of reaction of magnesium with solution P and solution Q. Explain your answer. [3]

2. Hydrochloric acid reacts with sodium hydroxide in a neutralisation reaction.

(a) Write a balanced symbol equation for this reaction. Include state symbols. [2]

(b) A student carries out a titration to find the concentration of a sodium hydroxide solution.

The student pipettes 25.0 cm³ of sodium hydroxide solution into a conical flask and adds a few drops of methyl orange indicator. The sodium hydroxide solution is then titrated with 0.100 mol/dm³ hydrochloric acid.

The results are shown in the table below.

Titration	Initial reading (cm³)	Final reading (cm³)	Volume used (cm³)
1	0.0	26.8	26.8
2	0.0	26.2	26.2
3	0.0	26.0	26.0

(i) Calculate the average volume of hydrochloric acid used. [1]

(ii) Calculate the number of moles of hydrochloric acid used in the titration. [2]

(iii) Use your answer from (ii) to calculate the concentration of the sodium hydroxide solution in mol/dm³. [2]

(iv) Calculate the concentration of the sodium hydroxide solution in g/dm³. (Mr of NaOH = 40) [2]

3. A student investigates the preparation of salts using different methods.

(a) The student wants to prepare zinc chloride crystals starting from zinc metal.

(i) Name a suitable acid to use. [1]

(ii) Describe the method the student should use to prepare pure, dry crystals of zinc chloride. [4]

(b) The student wants to prepare barium sulfate.

(i) Suggest two suitable starting materials for this preparation. [2]

(ii) Name the method used to prepare barium sulfate. [1]

(iii) Explain why this method is necessary for preparing barium sulfate. [2]

4. Ammonia is an important industrial chemical.

(a) Ammonia is manufactured by the Haber Process.

(i) Write a balanced symbol equation for the formation of ammonia from its elements. [2]

(ii) State two conditions used in the Haber Process. [2]

(b) Ammonia dissolves in water to form an alkaline solution.

(i) Write an equation to show how ammonia produces hydroxide ions in water. [1]

(ii) A student adds a few drops of ammonia solution to separate samples of the following metal salt solutions:

• Copper(II) sulfate
• Iron(II) sulfate
• Aluminium chloride

Complete the table below to show the observations made. [3]

Metal salt solution	Observation with ammonia solution
Copper(II) sulfate
Iron(II) sulfate
Aluminium chloride

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Section B [40 marks]

5. A student investigates the reaction between metals and acids.

The student adds 1.30 g of zinc powder to 50.0 cm³ of 1.00 mol/dm³ hydrochloric acid in a conical flask. The reaction produces zinc chloride and hydrogen gas.

(a) Write a balanced symbol equation for this reaction. Include state symbols. [2]

(b) Calculate the number of moles of zinc used. (Ar: Zn = 65) [2]

(c) Calculate the number of moles of hydrochloric acid used. [2]

(d) Determine which reactant is in excess and calculate how many moles are in excess. [3]

(e) Calculate the volume of hydrogen gas produced at room temperature and pressure. [2]

(f) The student repeats the experiment using the same mass of zinc but with 50.0 cm³ of 1.00 mol/dm³ ethanoic acid instead of hydrochloric acid.

(i) State two differences you would expect to observe compared to the original experiment. [2]

(ii) Explain these differences in terms of acid strength. [3]

6. A student investigates the thermal decomposition of metal carbonates.

The student heats samples of different metal carbonates and records the observations.

(a) Complete the word equation for the thermal decomposition of calcium carbonate: [1]

calcium carbonate → _____________ + _____________

(b) Write a balanced symbol equation for the thermal decomposition of copper(II) carbonate. [2]

(c) The student heats 2.48 g of copper(II) carbonate until no further change occurs. Calculate the mass of copper(II) oxide formed. (Mr: CuCO₃ = 124, CuO = 80) [3]

(d) The student notices that some metal carbonates decompose more easily than others.

(i) Suggest which carbonate would decompose at the lowest temperature: sodium carbonate, magnesium carbonate, or copper(II) carbonate. [1]

(ii) Explain your answer in terms of the position of the metals in the reactivity series. [2]

7. Paper chromatography is used to separate and identify compounds.

A student uses paper chromatography to analyse a sample of food colouring. The chromatogram obtained is shown below.

[THIS IS FIGURE: A chromatography diagram showing:

• Solvent front at 8.0 cm from start line
• Food colouring sample showing 2 spots: one at 2.4 cm and another at 6.4 cm from start line
• Reference compounds A, B, C, and D showing spots at various heights]

(a) Calculate the Rf value for each spot in the food colouring sample. [2]

Spot 1: Rf = _____________

Spot 2: Rf = _____________

(b) The reference compounds have the following Rf values:

• Compound A: 0.30
• Compound B: 0.80
• Compound C: 0.45
• Compound D: 0.60

Use this information to identify which compounds are present in the food colouring sample. [2]

(c) Explain why the compounds separate during chromatography. [2]

(d) Suggest how the student could improve the separation of the compounds. [1]

8. A student investigates electrolysis using different electrolytes.

(a) Define electrolysis. [2]

(b) The student electrolyses molten sodium chloride using inert electrodes.

(i) Name the products formed at each electrode. [2]

Cathode (negative electrode): _____________

Anode (positive electrode): _____________

(ii) Write ionic equations for the reactions at each electrode. [2]

Cathode: _____________

Anode: _____________

(c) The student then electrolyses dilute sodium chloride solution using inert electrodes.

(i) Explain why the products are different from those obtained with molten sodium chloride. [3]

(ii) Write an ionic equation for the reaction at the cathode when dilute sodium chloride solution is electrolysed. [1]

(d) Describe one industrial use of electrolysis. [2]

TuitionGoWhere Practice Paper - Chemistry O-Level - Mark Scheme

Total: 80 marks

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Section A [40 marks]

1. Acids and bases investigation

(a) Complete the statements: [2] (i) H⁺ (hydrogen) ions [1] (ii) OH⁻ (hydroxide) ions [1]

(b) Universal Indicator results: (i) Which solution is neutral? [1] R [1]

(ii) Which solution is the strongest acid? Explain. [2] P [1] - because it has the lowest pH/highest H⁺ ion concentration [1]

(iii) Arrange P, Q and S in order of increasing H⁺ concentration. [1] S, Q, P [1]

(c) Magnesium with acids: (i) Word equation: [1] magnesium + acid → salt + hydrogen [1]

(ii) Two observations with solution P: [2]

• Fizzing/effervescence/bubbles [1]
• Magnesium dissolves/disappears [1] Accept: Heat produced, squeaky pop with lighted splint

(iii) Compare rates with P and Q: [3] P reacts faster than Q [1] P is a stronger acid than Q [1]
P has higher H⁺ ion concentration, so more frequent collisions between H⁺ ions and magnesium [1]

2. Neutralisation and titration

(a) Balanced equation with state symbols: [2] HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l) [2] 1 mark for correct formulae and balancing, 1 mark for state symbols

(b) Titration calculations: (i) Average volume of HCl: [1] 26.1 cm³ [1] Working: (26.2 + 26.0) ÷ 2 = 26.1 cm³ (ignore first titration as rough)

(ii) Moles of HCl used: [2] n = c × V = 0.100 × (26.1/1000) = 0.00261 mol [2] 1 mark for method, 1 mark for answer

(iii) Concentration of NaOH: [2] From equation: 1 mol HCl : 1 mol NaOH Moles NaOH = 0.00261 mol c = n/V = 0.00261/(25.0/1000) = 0.104 mol/dm³ [2] 1 mark for stoichiometry, 1 mark for calculation

(iv) Concentration in g/dm³: [2] c = 0.104 × 40 = 4.16 g/dm³ [2] 1 mark for method, 1 mark for answer

3. Salt preparation methods

(a) Zinc chloride preparation: (i) Suitable acid: [1] Hydrochloric acid [1]

(ii) Method for pure, dry crystals: [4]

1. Add excess zinc to dilute hydrochloric acid [1]
2. Filter to remove unreacted zinc [1]
3. Evaporate the filtrate to concentrate the solution [1]
4. Cool to allow crystallisation, then filter and dry the crystals [1]

(b) Barium sulfate preparation: (i) Two suitable starting materials: [2] Barium chloride/nitrate and sodium/potassium sulfate [2] Accept any soluble barium salt + any soluble sulfate

(ii) Method name: [1] Precipitation [1]

(iii) Why this method is necessary: [2] Barium sulfate is insoluble in water [1] Cannot be prepared by evaporation methods [1]

4. Ammonia

(a) Haber Process: (i) Balanced equation: [2] N₂(g) + 3H₂(g) ⇌ 2NH₃(g) [2] 1 mark for correct formulae and balancing, 1 mark for reversible arrow

(ii) Two conditions: [2] High pressure (200-300 atm) [1] High temperature (450°C) OR Iron catalyst [1]

(b) Ammonia in water: (i) Equation for OH⁻ production: [1] NH₃(aq) + H₂O(l) → NH₄⁺(aq) + OH⁻(aq) [1]

(ii) Observations with metal salt solutions: [3]

Metal salt solution	Observation
Copper(II) sulfate	Blue precipitate [1]
Iron(II) sulfate	Green precipitate [1]
Aluminium chloride	White precipitate [1]

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Section B [40 marks]

5. Metals and acids investigation

(a) Balanced equation: [2] Zn(s) + 2HCl(aq) → ZnCl₂(aq) + H₂(g) [2]

(b) Moles of zinc: [2] n = m/Ar = 1.30/65 = 0.0200 mol [2]

(c) Moles of HCl: [2] n = c × V = 1.00 × (50.0/1000) = 0.0500 mol [2]

(d) Excess reactant calculation: [3] From equation: 1 mol Zn : 2 mol HCl 0.0200 mol Zn needs 0.0400 mol HCl [1] HCl is in excess [1] Excess HCl = 0.0500 - 0.0400 = 0.0100 mol [1]

(e) Volume of hydrogen gas: [2] From equation: 1 mol Zn produces 1 mol H₂ 0.0200 mol H₂ produced V = n × 24 = 0.0200 × 24 = 0.480 dm³ [2]

(f) Comparison with ethanoic acid: (i) Two differences: [2] Slower rate of reaction [1] Less vigorous fizzing/bubbling [1]

(ii) Explanation in terms of acid strength: [3] Ethanoic acid is a weak acid [1] Only partially ionises in water [1] Lower concentration of H⁺ ions, so fewer collisions with zinc [1]

6. Thermal decomposition of carbonates

(a) Word equation completion: [1] calcium carbonate → calcium oxide + carbon dioxide [1]

(b) Balanced equation for copper(II) carbonate: [2] CuCO₃(s) → CuO(s) + CO₂(g) [2]

(c) Mass of CuO formed: [3] Moles of CuCO₃ = 2.48/124 = 0.0200 mol [1] From equation: 1 mol CuCO₃ → 1 mol CuO [1] Mass of CuO = 0.0200 × 80 = 1.60 g [1]

(d) Thermal stability: (i) Carbonate decomposing at lowest temperature: [1] Copper(II) carbonate [1]

(ii) Explanation using reactivity series: [2] Copper is less reactive than sodium and magnesium [1] Less reactive metals have carbonates that decompose more easily [1]

7. Paper chromatography

(a) Rf value calculations: [2] Spot 1: Rf = 2.4/8.0 = 0.30 [1] Spot 2: Rf = 6.4/8.0 = 0.80 [1]

(b) Compound identification: [2] Spot 1 = Compound A (Rf = 0.30) [1] Spot 2 = Compound B (Rf = 0.80) [1]

(c) Why compounds separate: [2] Different compounds have different solubilities in the solvent [1] Different compounds have different attractions to the paper [1]

(d) Improve separation: [1] Use a different solvent OR Use a longer piece of chromatography paper [1]

8. Electrolysis

(a) Definition of electrolysis: [2] The breakdown/decomposition of an ionic compound [1] by passing an electric current through it when molten or in solution [1]

(b) Molten sodium chloride: (i) Products at electrodes: [2] Cathode: Sodium [1] Anode: Chlorine [1]

(ii) Ionic equations: [2] Cathode: Na⁺ + e⁻ → Na [1] Anode: 2Cl⁻ → Cl₂ + 2e⁻ [1]

(c) Dilute sodium chloride solution: (i) Why products are different: [3] Water is present in the solution [1] H⁺ ions from water are preferentially discharged at cathode instead of Na⁺ [1] Because hydrogen is less reactive than sodium [1]

(ii) Ionic equation at cathode: [1] 2H⁺ + 2e⁻ → H₂ [1]

(d) Industrial use of electrolysis: [2] Extraction of aluminium from aluminium oxide [1] Electroplating metals [1] Accept: Purification of copper, Production of chlorine/sodium hydroxide