Secondary 4 Pure Chemistry Practice Paper 2

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TuitionGoWhere Practice Paper - Pure Chemistry Secondary 4

TuitionGoWhere Practice Paper (AI) - Version 2

Subject: Pure Chemistry
Level: Secondary 4
Paper: Structured and Free Response
Duration: 1 hour 45 minutes
Total Marks: 80
Name: __________________________ Class: __________ Date: __________

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Instructions to Candidates

1. Answer all questions in the spaces provided.
2. Write your answers clearly and concisely.
3. State all chemical equations with balanced symbols and state symbols where required.
4. Use of a scientific calculator is permitted.

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Section A: Structured Questions (50 Marks)

Question 1 A student is provided with three colorless solutions: sodium chloride, sodium sulfate, and sodium carbonate. (a) Describe a chemical test to identify the sodium carbonate solution. [2]

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(b) The student then wishes to differentiate between the sodium chloride and sodium sulfate solutions. Suggest a suitable reagent and describe the observation for each. [3]

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Question 2 The reaction between dilute sulfuric acid and an alkali is studied. (a) Define the term neutralisation in terms of ions. [1]

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(b) A student reacts 25.0 $\text{cm}^3$ of 0.10 $\text{mol/dm}^3$ sodium hydroxide solution with excess dilute sulfuric acid. (i) Write the balanced chemical equation for the reaction, including state symbols. [2]

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(ii) Calculate the volume of 0.10 $\text{mol/dm}^3$ sulfuric acid required to exactly neutralize the sodium hydroxide. [3]

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Question 3 The Haber Process is used for the industrial manufacture of ammonia. (a) State the chemical equation for the production of ammonia. [2]

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(b) The reaction is reversible. Explain how an increase in pressure affects the yield of ammonia. [2]

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(c) Why is a catalyst used in this process, and what is the specific catalyst used? [2]

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Question 4 A sample of an impure metal carbonate is reacted with dilute hydrochloric acid. (a) State two observations made during this reaction. [2]

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(b) The gas evolved is passed through lime water. State the observation and explain why this confirms the gas is carbon dioxide. [2]

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(c) If the metal carbonate was calcium carbonate, write the balanced equation for the reaction with $\text{HCl}$. [2]

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Question 5 Solubility rules are essential for predicting the outcome of reactions. (a) State whether the following salts are soluble or insoluble in water: [3]

• Silver Nitrate: _________________
• Barium Sulfate: _________________
• Potassium Carbonate: _________________ (b) Describe a method to prepare a pure sample of lead(II) iodide. [3]

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Question 6 Strong and weak acids differ in their ionization. (a) Explain the difference between a strong acid and a weak acid in terms of the extent of ionization in aqueous solution. [2]

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(b) Compare the pH of 0.1 $\text{mol/dm}^3$ $\text{HCl}$ and 0.1 $\text{mol/dm}^3$ $\text{CH}_3\text{COOH}$. Explain your answer. [3]

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Question 7 A student investigates the reaction of different metals with dilute nitric acid. (a) Why is it generally unsuitable to use very reactive metals like potassium to prepare salts in a school laboratory? [2]

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(b) Magnesium reacts with dilute $\text{HCl}$ to produce a salt. State the name of the salt and the observation during the reaction. [2]

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Question 8 Amphoteric oxides can react with both acids and alkalis. (a) Give an example of an amphoteric oxide. [1]

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(b) Write a balanced equation showing how this oxide reacts with sodium hydroxide. [2]

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Question 9 The pH of soil affects crop growth. (a) If a soil is too acidic, what substance can a farmer add to the soil to neutralize it? [1]

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(b) Explain the chemical process that occurs when this substance is added to the soil. [2]

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Question 10 A salt is analyzed and found to contain the $\text{NO}_3^-$ ion. (a) Describe the test for the nitrate ion using aqueous sodium hydroxide and silver nitrate. [3]

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(b) State the color of the precipitate formed if the salt also contained $\text{Cu}^{2+}$ ions. [1]

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Section B: Free-Response Questions (30 Marks)

Question 11 A student is tasked with preparing a pure sample of Zinc Sulfate. (a) Suggest the most suitable starting materials for this preparation. [2]

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(b) Describe the full experimental procedure, from the initial reaction to obtaining dry crystals. [6]

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(c) Explain why the zinc powder is added in excess and how the excess is removed. [3]

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Question 12 Consider the reaction: $\text{NH}_3(\text{g}) + \text{HCl}(\text{g}) \rightarrow \text{NH}_4\text{Cl}(\text{s})$. (a) Describe the observation when a glass rod dipped in concentrated $\text{HCl}$ is brought near a jar of ammonia gas. [2]

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(b) Explain the bonding in the product $\text{NH}_4\text{Cl}$ in terms of its structure. [3]

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(c) Compare the solubility of $\text{NH}_4\text{Cl}$ with that of $\text{AgCl}$. Explain your answer. [3]

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Question 13 The effect of acid rain on the environment is significant. (a) Explain how sulfur dioxide ($\text{SO}_2$) is formed during the combustion of fossil fuels and how it leads to acid rain. [4]

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(b) Describe the effect of acid rain on limestone statues. Write the chemical equation for the reaction. [4]

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(c) Suggest one way to reduce the emission of $\text{SO}_2$ from power stations. [2]

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Answer Key - Pure Chemistry Secondary 4 Practice Paper (Version 2)

Section A: Structured Questions

Q1 (a) Add dilute $\text{HCl}$ to the solution. Observation: Effervescence / bubbles of colorless gas evolved which turns lime water milky. [2] (b) Reagent: Silver nitrate solution ($\text{AgNO}_3$).

• Sodium chloride: White precipitate formed. [1]
• Sodium sulfate: White precipitate formed. [1]
• Differentiation: Add dilute $\text{HNO}_3$ or observe that $\text{AgCl}$ is soluble in dilute ammonia while $\text{Ag}_2\text{SO}_4$ is not (or use $\text{BaCl}_2$ to get white ppt for sulfate and no ppt for chloride). [1]

Q2 (a) The reaction between an acid and a base to form a salt and water. / The reaction where $\text{H}^+$ ions from an acid combine with $\text{OH}^-$ ions from a base to form water. [1] (b) (i) $2\text{NaOH}(\text{aq}) + \text{H}_2\text{SO}_4(\text{aq}) \rightarrow \text{Na}_2\text{SO}_4(\text{aq}) + 2\text{H}_2\text{O}(\text{l})$ [2] (ii) Moles $\text{NaOH} = 0.025 \times 0.10 = 0.0025 \text{ mol}$. Moles $\text{H}_2\text{SO}_4 = 0.0025 / 2 = 0.00125 \text{ mol}$. Volume $\text{H}_2\text{SO}_4 = 0.00125 / 0.10 = 0.0125 \text{ dm}^3 = 12.5 \text{ cm}^3$. [3]

Q3 (a) $\text{N}(\text{g}) + 3\text{H}_2(\text{g}) \rightleftharpoons 2\text{NH}_3(\text{g})$ [2] (b) Increasing pressure shifts the equilibrium to the right (towards the side with fewer gas molecules), increasing the yield of ammonia. [2] (c) Catalyst: Iron (Fe). Purpose: To increase the rate of reaction by lowering the activation energy. [2]

Q4 (a) Effervescence / bubbles of gas; solution becomes colorless (if carbonate was colored). [2] (b) Observation: Lime water turns milky/cloudy. Explanation: $\text{CO}_2$ reacts with $\text{Ca}(\text{OH})_2$ to form insoluble $\text{CaCO}_3$. [2] (c) $\text{CaCO}_3(\text{s}) + 2\text{HCl}(\text{aq}) \rightarrow \text{CaCl}_2(\text{aq}) + \text{H}_2\text{O}(\text{l}) + \text{CO}_2(\text{g})$ [2]

Q5 (a) Silver Nitrate: Soluble; Barium Sulfate: Insoluble; Potassium Carbonate: Soluble. [3] (b) Precipitation method. Mix aqueous lead(II) nitrate and aqueous potassium iodide. Filter the yellow precipitate of lead(II) iodide, wash with distilled water, and dry. [3]

Q6 (a) Strong acids ionize completely in aqueous solution to produce a high concentration of $\text{H}^+$ ions. Weak acids ionize only partially, producing a low concentration of $\text{H}^+$ ions. [2] (b) $\text{HCl}$ has a lower pH (more acidic). [1] $\text{HCl}$ is a strong acid that fully dissociates, whereas $\text{CH}_3\text{COOH}$ is a weak acid that only partially dissociates, resulting in a lower concentration of $\text{H}^+$ ions for the same molarity. [2]

Q7 (a) The reaction is too vigorous/violent and may cause the reaction mixture to spray or explode, making it dangerous and difficult to control. [2] (b) Salt: Magnesium chloride. Observation: Vigorous effervescence / bubbles of colorless gas. [2]

Q8 (a) Aluminum oxide ($\text{Al}_2\text{O}_3$) or Zinc oxide ($\text{ZnO}$). [1] (b) $\text{Al}_2\text{O}_3(\text{s}) + 2\text{NaOH}(\text{aq}) + 3\text{H}_2\text{O}(\text{l}) \rightarrow 2\text{NaAl}(\text{OH})_4(\text{aq})$ (or simplified version). [2]

Q9 (a) Slaked lime / Calcium hydroxide / Calcium carbonate. [1] (b) The basic substance reacts with the $\text{H}^+$ ions in the soil (neutralization), increasing the pH to a level suitable for plant growth. [2]

Q10 (a) Add $\text{NaOH}(\text{aq})$ to warm the solution, then add $\text{AgNO}_3(\text{aq})$. A brown ring is formed (if using the brown ring test) or simply describe the specific nitrate test sequence. [3] (b) Blue precipitate. [1]

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Section B: Free-Response Questions

Q11 (a) Zinc powder and dilute sulfuric acid. [2] (b) 1. Add excess zinc powder to warm dilute sulfuric acid. [1] 2. Stir until no more gas is evolved. [1] 3. Filter the mixture to remove unreacted zinc. [1] 4. Heat the filtrate (zinc sulfate solution) in an evaporating dish until the crystallization point is reached. [1] 5. Allow to cool slowly to form crystals. [1] 6. Filter crystals and pat dry with filter paper. [1] (c) To ensure all the sulfuric acid has reacted completely. [2] Removed by filtration. [1]

Q12 (a) Dense white fumes/smoke are formed. [2] (b) $\text{NH}_4\text{Cl}$ is an ionic compound. It consists of ammonium ions ($\text{NH}_4^+$) and chloride ions ($\text{Cl}^-$) held together by strong electrostatic forces of attraction in a giant lattice. [3] (c) $\text{NH}_4\text{Cl}$ is soluble; $\text{AgCl}$ is insoluble. [1] Ammonium salts are generally soluble in water, whereas silver halides (except $\text{AgF}$) are insoluble. [2]

Q13 (a) Sulfur impurities in fossil fuels react with oxygen during combustion to form $\text{SO}_2$. [2] $\text{SO}_2$ is further oxidized to $\text{SO}_3$, which reacts with water vapor in the atmosphere to form sulfuric acid ($\text{H}_2\text{SO}_4$), resulting in acid rain. [2] (b) Acid rain reacts with calcium carbonate in limestone to form soluble calcium salts and $\text{CO}_2$, causing the stone to erode/corrode. [2] $\text{CaCO}_3(\text{s}) + \text{H}_2\text{SO}_4(\text{aq}) \rightarrow \text{CaSO}_4(\text{s}) + \text{H}_2\text{O}(\text{l}) + \text{CO}_2(\text{g})$ [2] (c) Use of flue-gas desulfurization (scrubbing with $\text{CaCO}_3$ or $\text{CaO}$) to remove $\text{SO}_2$ before it leaves the chimney. [2]