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Secondary 4 Pure Chemistry Practice Paper 2

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TuitionGoWhere Practice Paper - Pure Chemistry Secondary 4

TuitionGoWhere Practice Paper (AI) Version: 2 of 5 Subject: Pure Chemistry (6092) Level: Secondary 4 Paper: Practice Paper – Topic: Acids, Bases and Salts Duration: 1 hour 15 minutes Total Marks: 50

Name: __________________________
Class: __________________________
Date: __________________________

Instructions to Candidates

Write your name, class, and date in the spaces provided.
Answer all questions.
Write your answers in the spaces provided on the question paper.
The number of marks is given in brackets [ ] at the end of each question or part question.
A copy of the Periodic Table is printed on page 12 (not included in this extract, assume standard access).
You may use a calculator.

Section A: Structured Questions (30 Marks)

Answer all questions in this section.

1. Dilute sulfuric acid reacts with excess copper(II) carbonate to form copper(II) sulfate, water, and carbon dioxide.

(a) Describe the observations you would make during this reaction.
........................................................................................................................................
........................................................................................................................................ [2]

(b) Write a balanced chemical equation for this reaction, including state symbols.
........................................................................................................................................ [2]

(c) Explain why copper metal cannot be used instead of copper(II) carbonate to prepare copper(II) sulfate using dilute sulfuric acid.
........................................................................................................................................
........................................................................................................................................ [1]

2. A student investigates the reaction between aqueous sodium hydroxide and dilute hydrochloric acid.

(a) Name the type of chemical reaction that occurs.
........................................................................................................................................ [1]

(b) The student adds universal indicator to the hydrochloric acid and then adds sodium hydroxide dropwise until the solution is neutral. State the colour change observed.
From: __________________________ To: __________________________ [1]

(c) Write the ionic equation for this neutralisation reaction.
........................................................................................................................................ [1]

3. Salt X is a white solid. When heated strongly, it decomposes to form a yellow solid that turns white on cooling, and a brown gas is evolved.

(a) Identify the brown gas.
........................................................................................................................................ [1]

(b) Identify the cation present in Salt X.
........................................................................................................................................ [1]

(c) Salt X is soluble in water. Describe a chemical test to confirm the presence of the anion in Salt X.
Test: ...........................................................................................................................
Observation: ............................................................................................................... [2]

4. The table below shows the pH values of four different aqueous solutions, P, Q, R, and S.

Solution	pH
P	1
Q	7
R	13
S	9

(a) Which solution has the highest concentration of hydrogen ions, $H^+$ ?
........................................................................................................................................ [1]

(b) Which solution could be aqueous ammonia?
........................................................................................................................................ [1]

(c) Solution P is a strong acid. Explain, in terms of ionisation, what is meant by a "strong acid".
........................................................................................................................................
........................................................................................................................................ [1]

5. Zinc oxide is an amphoteric oxide.

(a) Define the term amphoteric oxide.
........................................................................................................................................
........................................................................................................................................ [1]

(b) Write balanced chemical equations for the reaction of zinc oxide with:
(i) Dilute hydrochloric acid.
........................................................................................................................................ [1]
(ii) Aqueous sodium hydroxide.
........................................................................................................................................ [1]

6. Barium chloride solution is added to solution Y. A white precipitate forms which is insoluble in dilute nitric acid.

(a) Identify the anion present in solution Y.
........................................................................................................................................ [1]

(b) Why is dilute nitric acid added before testing for this anion?
........................................................................................................................................
........................................................................................................................................ [1]

(c) If dilute hydrochloric acid were used instead of dilute nitric acid, explain why the test result might be invalid.
........................................................................................................................................
........................................................................................................................................ [1]

7. Ammonia gas is manufactured by the Haber Process.

$N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g) \quad \Delta H = -92 \text{ kJ/mol}$

(a) State the catalyst used in the Haber Process.
........................................................................................................................................ [1]

(b) Explain why a high pressure is used in the Haber Process, referring to both yield and rate.
........................................................................................................................................
........................................................................................................................................
........................................................................................................................................ [2]

8. A student wants to prepare pure, dry crystals of magnesium sulfate ( $MgSO_4$ ) from magnesium carbonate and dilute sulfuric acid.

(a) Why is an excess of magnesium carbonate used?
........................................................................................................................................ [1]

(b) After the reaction is complete, the mixture is filtered. What is removed by filtration?
........................................................................................................................................ [1]

(c) Describe the subsequent steps to obtain pure, dry crystals from the filtrate.
........................................................................................................................................
........................................................................................................................................
........................................................................................................................................ [2]

Section B: Free-Response Questions (20 Marks)

Answer all questions in this section.

9. Hydrochloric acid reacts with calcium carbonate according to the following equation:

$CaCO_3(s) + 2HCl(aq) \rightarrow CaCl_2(aq) + H_2O(l) + CO_2(g)$

In an experiment, excess calcium carbonate is added to $50.0 \text{ cm}^3$ of $2.0 \text{ mol/dm}^3$ hydrochloric acid. The volume of carbon dioxide gas collected is measured every minute.

(a) Calculate the maximum volume of carbon dioxide gas produced at room temperature and pressure (r.t.p.).
[Molar volume of gas at r.t.p. = $24 \text{ dm}^3$ ]
[3]

(b) The experiment is repeated using $50.0 \text{ cm}^3$ of $1.0 \text{ mol/dm}^3$ hydrochloric acid. All other conditions remain the same.
On the grid below, sketch the curve for the original experiment (label it A) and the new experiment (label it B).

Image pending generation for this question.

[2]

(c) Explain, in terms of collision theory, why the initial rate of reaction in experiment B is slower than in experiment A.
........................................................................................................................................
........................................................................................................................................
........................................................................................................................................ [2]

10. Solution A contains iron(II) ions ( $Fe^{2+}$ ). Solution B contains iron(III) ions ( $Fe^{3+}$ ).

(a) Describe the observations when aqueous sodium hydroxide is added separately to:
(i) Solution A
........................................................................................................................................ [1]
(ii) Solution B
........................................................................................................................................ [1]

(b) Both precipitates formed in (a) are left standing in air for a period of time. Describe and explain any change in appearance of the precipitate from Solution A.
........................................................................................................................................
........................................................................................................................................
........................................................................................................................................ [2]

(c) Iron(III) oxide can be reduced to iron using carbon monoxide in a blast furnace.
(i) Write the balanced chemical equation for this reaction.
........................................................................................................................................ [1]
(ii) Identify the reducing agent in this reaction and explain your choice in terms of oxygen transfer.
........................................................................................................................................
........................................................................................................................................ [2]

11. Ethanoic acid ( $CH_3COOH$ ) is a weak organic acid, while hydrochloric acid ( $HCl$ ) is a strong mineral acid. Both acids react with magnesium ribbon.

(a) Explain why ethanoic acid is classified as a weak acid.
........................................................................................................................................
........................................................................................................................................ [1]

(b) When magnesium ribbon is added to separate solutions of $1.0 \text{ mol/dm}^3$ ethanoic acid and $1.0 \text{ mol/dm}^3$ hydrochloric acid, the reaction with hydrochloric acid is much more vigorous.
Explain this difference in reaction rate.
........................................................................................................................................
........................................................................................................................................
........................................................................................................................................ [2]

(c) Despite the difference in rate, both acids will neutralise the same amount of sodium hydroxide if equal volumes and concentrations are used. Explain why.
........................................................................................................................................
........................................................................................................................................ [1]

12. Potassium nitrate ( $KNO_3$ ) is a soluble salt. It can be prepared by titration.

(a) Name the suitable acid and alkali required to prepare potassium nitrate.
Acid: __________________________
Alkali: __________________________ [1]

(b) Describe how you would use titration to prepare a pure sample of potassium nitrate solution. Include the use of an indicator in your answer.
........................................................................................................................................
........................................................................................................................................
........................................................................................................................................
........................................................................................................................................
........................................................................................................................................ [3]

(c) Why is the titration method preferred over the "excess solid" method for preparing potassium nitrate?
........................................................................................................................................
........................................................................................................................................ [1]

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Answers

TuitionGoWhere Practice Paper - Pure Chemistry Secondary 4

Answer Key & Marking Scheme

Version: 2 of 5 Topic: Acids, Bases and Salts

Section A: Structured Questions

1.
(a) Observations:

Effervescence / Bubbles of gas produced. [1]
Blue solid (copper(II) carbonate) dissolves / disappears to form a blue solution. [1]
(Note: "Solid disappears" alone is insufficient; must mention colour or gas.)

(b) Equation:
$CuCO_3(s) + H_2SO_4(aq) \rightarrow CuSO_4(aq) + H_2O(l) + CO_2(g)$
[1 for correct formulae, 1 for balancing and state symbols]

(c) Explanation:
Copper is below hydrogen in the reactivity series. [1]
(Therefore, it does not react with dilute acids to displace hydrogen.)

2.
(a) Neutralisation. [1]

(b) From: Red (or Orange/Pink depending on initial concentration, but Red is standard for strong acid) [0.5]
To: Green [0.5]
(Accept "Red to Green" or "Pink to Green" if methyl orange/universal indicator logic is applied correctly for neutral point. Universal Indicator is Green at pH 7.)

(c) Ionic Equation:
$H^+(aq) + OH^-(aq) \rightarrow H_2O(l)$
[1 for correct ions and product, state symbols optional but recommended]

3.
(a) Nitrogen dioxide ( $NO_2$ ). [1]

(b) Lead(II) ion ( $Pb^{2+}$ ). [1]
(Lead(II) nitrate decomposes to Lead(II) oxide (yellow when hot, white when cold) and $NO_2$ .)

(c) Test for Nitrate ( $NO_3^-$ ):
Test: Add aqueous sodium hydroxide and aluminium foil (or Devarda's alloy) and warm gently. [1]
Observation: Gas evolved that turns damp red litmus paper blue. [1]
(Alternative: Brown ring test is also acceptable but less common at this level for simple identification.)

4.
(a) Solution P. [1]
(Lower pH = higher $[H^+]$ .)

(b) Solution S. [1]
(Aqueous ammonia is a weak alkali, typically pH 8-10. Solution R (pH 13) is a strong alkali like NaOH.)

5.
(a) Definition:
An amphoteric oxide is an oxide that reacts with both acids and bases to form salt and water. [1]

(b) Equations:
(i) $ZnO(s) + 2HCl(aq) \rightarrow ZnCl_2(aq) + H_2O(l)$ [1]
(ii) $ZnO(s) + 2NaOH(aq) \rightarrow Na_2ZnO_2(aq) + H_2O(l)$ [1]
(Accept $Na_2[Zn(OH)_4]$ for the complex ion product if taught, but sodium zincate is standard.)

6.
(a) Sulfate ion ( $SO_4^{2-}$ ). [1]

(b) Reason:
To remove carbonate ions ( $CO_3^{2-}$ ) or sulfite ions ( $SO_3^{2-}$ ) which also form white precipitates with barium ions but dissolve in acid. [1]
(Prevents false positive.)

(c) Explanation:
Hydrochloric acid contains chloride ions ( $Cl^-$ ). If the solution contained silver ions ( $Ag^+$ ), a white precipitate of $AgCl$ would form, interfering with the test. More importantly, if testing for sulfate, adding HCl introduces no sulfate, but standard protocol uses nitric acid to avoid introducing anions that might precipitate with other potential cations or interfere. Specifically, if the unknown contained Lead(II), $PbCl_2$ is sparingly soluble and might confuse results, whereas nitrates are all soluble. [1]
(Key point: Avoid introducing the anion being tested for or interfering ions.)

7.
(a) Iron (Fe). [1]

(b) Explanation:

Yield: High pressure favours the forward reaction because there are fewer moles of gas on the product side (2 moles) than on the reactant side (4 moles). This increases the yield of ammonia. [1]
Rate: High pressure increases the concentration of gas particles, leading to more frequent collisions and a faster rate of reaction. [1]

8.
(a) To ensure all the sulfuric acid reacts / is neutralised. [1]

(b) Unreacted / Excess magnesium carbonate. [1]

Heat the filtrate to evaporate some water / until saturated. [1]
Allow the solution to cool to crystallise. [1]
(Filter, wash with cold water, and dry between filter papers are implied final steps, but crystallisation is the key marking point.)

Section B: Free-Response Questions

9.
(a) Calculation:

Moles of $HCl = \text{Concentration} \times \text{Volume}$
$= 2.0 \times \frac{50}{1000} = 0.10 \text{ mol}$ [1]
From equation: 2 mol $HCl$ produces 1 mol $CO_2$ .
Moles of $CO_2 = \frac{0.10}{2} = 0.05 \text{ mol}$ [1]
Volume of $CO_2 = \text{Moles} \times 24 \text{ dm}^3$
$= 0.05 \times 24 = 1.2 \text{ dm}^3$ [1]

(b) Graph:

Curve A: Starts steep, levels off at $1.2 \text{ dm}^3$ . [1]
Curve B: Starts less steep (slower rate), levels off at $0.6 \text{ dm}^3$ (half the moles of acid). [1]
(Check: B must be lower final volume and slower initial gradient.)

Solution B has a lower concentration of $H^+$ ions. [1]
This leads to fewer frequent collisions between $H^+$ ions and $CaCO_3$ particles per unit time. [1]
(Note: Must mention frequency of collisions. "Less energy" is incorrect for concentration changes.)

10.
(a) Observations:
(i) Solution A ( $Fe^{2+}$ ): Dirty green precipitate formed. [1]
(ii) Solution B ( $Fe^{3+}$ ): Reddish-brown / Orange-brown precipitate formed. [1]

(b) Change in Appearance:
The green precipitate turns brown / reddish-brown. [1]
Explanation: Iron(II) hydroxide is oxidised by oxygen in the air to form Iron(III) hydroxide. [1]

(c) Reduction:
(i) $Fe_2O_3 + 3CO \rightarrow 2Fe + 3CO_2$ [1]
(ii) Reducing Agent: Carbon monoxide ( $CO$ ). [1]
Explanation: It removes oxygen from iron(III) oxide / It gains oxygen to form carbon dioxide. [1]

11.
(a) Explanation:
Ethanoic acid only partially ionises / dissociates in water. [1]
(Most molecules remain as $CH_3COOH$ .)

(b) Rate Explanation:

Hydrochloric acid is a strong acid and has a higher concentration of $H^+$ ions compared to ethanoic acid of the same concentration. [1]
Higher $[H^+]$ leads to more frequent effective collisions with magnesium atoms. [1]

(c) Neutralisation Capacity:
Both acids have the same number of moles of acid molecules initially. [1]
As the reaction proceeds, the equilibrium for ethanoic acid shifts to the right, releasing more $H^+$ ions until all acid molecules have reacted.
(Accept: "Both are monoprotic acids and have the same molar amount of potential $H^+$ ".)

12.
(a) Acid: Nitric acid ( $HNO_3$ ). [0.5]
Alkali: Potassium hydroxide ( $KOH$ ). [0.5]

(b) Titration Method:

Pipette a known volume of potassium hydroxide into a conical flask and add a few drops of indicator (e.g., phenolphthalein). [1]
Fill a burette with nitric acid and add it to the flask until the indicator changes colour (endpoint). Record the volume. [1]
Repeat the titration without the indicator using the exact same volumes to obtain a pure salt solution (free from indicator contamination). [1]

(c) Reason:
Both potassium hydroxide and nitric acid are soluble / There is no insoluble reactant to filter off. [1]
(The "excess solid" method requires an insoluble base/carbonate/metal to filter off the excess. Since both reactants here are soluble, you cannot separate them by filtration if one is in excess. Titration ensures exact neutralisation.)