Secondary 4 Pure Chemistry Preliminary Examination Paper 2

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TuitionGoWhere Exam Practice (AI)

Subject: Pure Chemistry
Level: Secondary 4
Paper: Preliminary Examination (Version 2 of 5)
Duration: 1 hour 45 minutes
Total Marks: 80

Name: ___________________________ Class: ___________ Date: ___________

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Instructions to Candidates

1. Write your name and details in the spaces provided.
2. Answer all questions in Section A and Section B.
3. Write your answers in the spaces provided.
4. Show all working clearly for calculation questions.
5. Use the following atomic masses: H=1, C=12, N=14, O=16, Na=23, Mg=24, Al=27, S=32, Cl=35.5, K=39, Ca=40, Fe=56, Cu=64.

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Section A: Structured Questions (50 Marks)

Question 1 A student is investigating the properties of an unknown gas, Gas X. Gas X is a colourless gas that turns damp blue litmus paper red and reacts with sodium hydroxide solution to produce a salt and water. (a) Identify Gas X. [1]

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(b) Write a balanced chemical equation, including state symbols, for the reaction between Gas X and sodium hydroxide. [2]

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(c) Explain why damp litmus paper must be used to test for the acidity of Gas X. [2]

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Question 2 A sample of a metal, Metal M, is reacted with dilute hydrochloric acid. (a) If Metal M is potassium, discuss the suitability of this reaction for preparing a salt in a school laboratory. [2]

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(b) State two observations that would be seen if Metal M is magnesium. [2]

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(c) Write the ionic equation for the reaction between magnesium and dilute hydrochloric acid. [2]

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Question 3 Two colorless solutions, Solution A (containing $\text{Al}^{3+}$ ions) and Solution B (containing $\text{Pb}^{2+}$ ions), are provided. (a) Describe a chemical test using aqueous sodium hydroxide to differentiate between Solution A and Solution B. [3]

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(b) State the observation for Solution A when excess sodium hydroxide is added. [1]

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(c) Write the balanced equation for the formation of the precipitate in Solution B. [2]

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Question 4 The Haber Process is used for the industrial manufacture of ammonia. (a) State the chemical equation for the manufacture of ammonia. [2]

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(b) State the catalyst and the typical temperature and pressure used in this process. [3]

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(c) Explain why a compromise temperature is used rather than a very low temperature, despite the reaction being exothermic. [3]

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Question 5 A student prepares a sample of barium sulfate. (a) State the method used to prepare this salt. [1]

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(b) Name the two soluble salts that must be reacted to obtain this product. [2]

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(c) Explain why the product must be filtered and washed with distilled water. [2]

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Question 6 Consider the following acids: Ethanoic acid ($\text{CH}_3\text{COOH}$) and Hydrochloric acid ($\text{HCl}$). (a) Classify each acid as either "strong" or "weak". [2]

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(b) Explain the difference in their classification in terms of ionisation in aqueous solution. [3]

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(c) Which acid would have a lower pH if both were provided at a concentration of $0.1\text{ mol dm}^{-3}$? Explain your answer. [3]

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Question 7 A carbonate salt, Salt Y, is heated strongly in a test tube. (a) State the observation made during the heating of Salt Y. [2]

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(b) A gas is evolved. Describe a test to confirm the identity of this gas. [2]

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(c) If Salt Y is calcium carbonate, write the balanced chemical equation for its thermal decomposition. [2]

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Question 8 A titration is carried out between $25.0\text{ cm}^3$ of $0.100\text{ mol dm}^{-3}$ sodium hydroxide and a solution of sulfuric acid. (a) Define the term "neutralisation". [2]

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(b) Calculate the number of moles of $\text{NaOH}$ used in the titration. [2]

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(c) If $20.0\text{ cm}^3$ of sulfuric acid was required to reach the endpoint, calculate the concentration of the sulfuric acid in $\text{mol dm}^{-3}$. [4]

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Section B: Free-Response Questions (30 Marks)

Question 9 (a) Sulfur is burned in air to form sulfur dioxide. Sulfur dioxide then reacts further to form a gas that contributes to acid rain. (i) Write the equation for the formation of sulfur dioxide. [2]

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(ii) Name the gas formed when sulfur dioxide is further oxidised. [1]

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(iii) Describe how this gas leads to the formation of acid rain. [2]

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(b) Discuss the use of calcium carbonate in the treatment of flue gases from power stations to reduce air pollution. [4]

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Question 10 (a) Compare the solubility of nitrates, chlorides, and sulfates. [3]

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(b) A student wishes to prepare a pure, dry sample of copper(II) sulfate crystals. (i) Suggest a suitable starting material other than copper(II) oxide. [1]

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(ii) Describe the steps required to obtain the crystals from the reaction mixture, including how to ensure the crystals are not dissolved during drying. [5]

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Question 11 (a) Explain the role of pH control in agriculture, specifically regarding the use of lime (calcium oxide) on acidic soils. [4]

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(b) Describe the reaction between an acid and a metal carbonate. Include the general word equation and the observations expected. [4]

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Answer Key - Pure Chemistry Preliminary (Version 2)

Section A: Structured Questions

Question 1 (a) Carbon dioxide ($\text{CO}_2$) [1] (b) $\text{CO}_2(\text{g}) + 2\text{NaOH}(\text{aq}) \rightarrow \text{Na}_2\text{CO}_3(\text{aq}) + \text{H}_2\text{O}(\text{l})$ [2] (c) $\text{CO}_2$ is a molecular gas and only exhibits acidic properties when dissolved in water to form carbonic acid ($\text{H}_2\text{CO}_3$). Without water, $\text{H}^+$ ions cannot be released to change the color of the litmus. [2]

Question 2 (a) Unsuitable. Potassium is extremely reactive; the reaction with dilute $\text{HCl}$ would be too violent/explosive and dangerous for a school lab. [2] (b) Vigorous effervescence (bubbles of gas); the test tube becomes hot (exothermic). [2] (c) $\text{Mg}(\text{s}) + 2\text{H}^+(\text{aq}) \rightarrow \text{Mg}^{2+}(\text{aq}) + \text{H}_2(\text{g})$ [2]

Question 3 (a) Add aqueous $\text{NaOH}$ dropwise to both solutions. Both will form white precipitates. Add excess $\text{NaOH}$. The precipitate in Solution A ($\text{Al}^{3+}$) will redissolve to form a colorless solution, while the precipitate in Solution B ($\text{Pb}^{2+}$) will not redissolve (or only partially). [3] (b) White precipitate dissolves in excess / colorless solution formed. [1] (c) $\text{Pb}^{2+}(\text{aq}) + 2\text{OH}^-(\text{aq}) \rightarrow \text{Pb}(\text{OH})_2(\text{s})$ [2]

Question 4 (a) $\text{N}_2(\text{g}) + 3\text{H}_2(\text{g}) \rightleftharpoons 2\text{NH}_3(\text{g})$ [2] (b) Catalyst: Iron (Fe). Temperature: $450^\circ\text{C}$. Pressure: $200\text{ atm}$. [3] (c) The forward reaction is exothermic. A low temperature would shift equilibrium to the right (increasing yield), but the rate of reaction would be too slow to be economically viable. A compromise temperature ensures a reasonable rate and yield. [3]

Question 5 (a) Precipitation. [1] (b) Barium chloride ($\text{BaCl}_2$) and Sodium sulfate ($\text{Na}_2\text{SO}_4$) (or any soluble barium/sulfate salts). [2] (c) Filtered to remove the insoluble barium sulfate from the solution. Washed with distilled water to remove any remaining soluble impurities (e.g., $\text{NaCl}$). [2]

Question 6 (a) Ethanoic acid: Weak; Hydrochloric acid: Strong. [2] (b) $\text{HCl}$ is fully ionised in aqueous solution (all molecules split into $\text{H}^+$ and $\text{Cl}^-$). Ethanoic acid is only partially ionised, meaning most molecules remain intact. [3] (c) Hydrochloric acid. Because it is a strong acid, it produces a higher concentration of $\text{H}^+$ ions in solution, resulting in a lower pH. [3]

Question 7 (a) Salt Y decomposes; a gas is evolved (bubbles/effervescence). [2] (b) Bubble the gas through limewater. The limewater will turn cloudy/milky. [2] (c) $\text{CaCO}_3(\text{s}) \rightarrow \text{CaO}(\text{s}) + \text{CO}_2(\text{g})$ [2]

Question 8 (a) A reaction between an acid and a base to produce a salt and water. [2] (b) $\text{moles} = \text{concentration} \times \text{volume} = 0.100 \times (25/1000) = 0.0025\text{ mol}$ [2] (c) $\text{NaOH} + \text{H}_2\text{SO}_4 \rightarrow \text{Na}_2\text{SO}_4 + 2\text{H}_2\text{O}$ Ratio $\text{NaOH}:\text{H}_2\text{SO}_4 = 2:1$ $\text{Moles of } \text{H}_2\text{SO}_4 = 0.0025 / 2 = 0.00125\text{ mol}$ $\text{Concentration} = 0.00125 / (20/1000) = 0.0625\text{ mol dm}^{-3}$ [4]

Section B: Free-Response Questions

Question 9 (a) (i) $\text{S}(\text{s}) + \text{O}_2(\text{g}) \rightarrow \text{SO}_2(\text{g})$ [2] (ii) Sulfur trioxide ($\text{SO}_3$) [1] (iii) $\text{SO}_3$ dissolves in rainwater to form sulfuric acid ($\text{H}_2\text{SO}_4$), which lowers the pH of the rain. [2] (b) $\text{CaCO}_3$ is used to react with $\text{SO}_2$ in the flue gas. This neutralises the acidic $\text{SO}_2$ gas, converting it into solid calcium sulfite ($\text{CaSO}_3$) and $\text{CO}_2$, thereby preventing $\text{SO}_2$ from entering the atmosphere and forming acid rain. [4]

Question 10 (a) All nitrates are soluble. Most chlorides are soluble (except $\text{AgCl}$, $\text{PbCl}_2$). Most sulfates are soluble (except $\text{BaSO}_4$, $\text{PbSO}_4$, $\text{CaSO}_4$). [3] (b) (i) Copper(II) carbonate ($\text{CuCO}_3$). [1] (ii) 1. Heat the solution to saturation. 2. Allow to cool slowly to form crystals. 3. Filter the crystals. 4. Pat dry with filter paper (do not heat strongly to avoid dehydration/decomposition). [5]

Question 11 (a) Acidic soils can inhibit plant growth. Lime ($\text{CaO}$) is a basic oxide that reacts with the acid in the soil (neutralisation), increasing the pH to a level suitable for the specific crop. [4] (b) Reaction: Acid + Metal Carbonate $\rightarrow$ Salt + Water + Carbon Dioxide. Observation: Effervescence/bubbling as $\text{CO}_2$ gas is released; the solid carbonate dissolves. [4]