A Level H2 Chemistry Practice Paper 4

TuitionGoWhere Practice Paper — Chemistry H2 A-Level

Answer Key: Acids, Bases & Salts (Version 4)

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Section A

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Question 1 [2 marks]

Answer: A strong acid is an acid that completely dissociates (or ionises) in aqueous solution to produce $H^+(aq)$ ions.

Marking notes:

• [1] for "completely dissociates/ionises"
• [1] for reference to aqueous solution and production of $H^+$ ions

Common mistakes:

• Saying "partially dissociates" — this describes a weak acid.
• Omitting "completely" — the key distinction between strong and weak acids is the extent of dissociation.

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Question 2 [3 marks]

(a) [1 mark]

Answer: $[H^+] = 10^{-pH} = 10^{-1.20} = 0.0631$ mol dm⁻³ (to 3 s.f.)

Working: $[H^+] = 10^{-1.20} = 0.0631 \text{ mol dm}^{-3}$

(b) [2 marks]

Answer: HCl is a strong monoprotic acid, so it dissociates completely: $HCl(aq) \rightarrow H^+(aq) + Cl^-(aq)$

Since each molecule of HCl produces one $H^+$ ion: $[HCl] = [H^+] = 0.0631 \text{ mol dm}^{-3}$

Marking notes:

• [1] for stating HCl is a strong acid / dissociates completely
• [1] for correct concentration = 0.0631 mol dm⁻³

Common mistakes:

• Forgetting that HCl is monoprotic (1:1 ratio of HCl to H⁺).
• Using $[H^+] = 10^{pH}$ instead of $10^{-pH}$.

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Question 3 [3 marks]

(a) [2 marks]

Answer: For a weak acid: $K_a = \frac{[H^+][A^-]}{[HA]}$

Since $[H^+] = [A^-]$ and $[HA] \approx 0.100$ mol dm⁻³ (initial concentration, as dissociation is small):

$K_a = \frac{[H^+]^2}{[HA]}$ $[H^+]^2 = K_a \times [HA] = 1.74 \times 10^{-5} \times 0.100 = 1.74 \times 10^{-6}$ $[H^+] = \sqrt{1.74 \times 10^{-6}} = 1.32 \times 10^{-3} \text{ mol dm}^{-3}$ $pH = -\log(1.32 \times 10^{-3}) = 2.88$

Marking notes:

• [1] for correct setup of $K_a$ expression and substitution
• [1] for correct pH = 2.88

(b) [1 mark]

Answer: The pH at the equivalence point is greater than 7 (approximately 8–9). This is because the salt formed, sodium ethanoate (CH₃COONa), is the salt of a weak acid and strong base. The ethanoate ion hydrolyses in water to produce $OH^-$ ions, making the solution slightly alkaline.

Common mistakes:

• Saying pH = 7 — this is only true for strong acid–strong base titrations.
• Not explaining the hydrolysis of the conjugate base.

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Question 4 [4 marks]

(a) [1 mark]

Answer: Increasing acid strength: HCN < CH₃COOH < HCOOH

(The larger the $K_a$, the stronger the acid.)

(b) [2 marks]

Answer: $K_a = \frac{[H^+]^2}{[HA]}$ $[H^+]^2 = 1.60 \times 10^{-4} \times 0.200 = 3.20 \times 10^{-5}$ $[H^+] = \sqrt{3.20 \times 10^{-5}} = 5.66 \times 10^{-3} \text{ mol dm}^{-3}$ $pH = -\log(5.66 \times 10^{-3}) = 2.25$

Marking notes:

• [1] for correct substitution into $K_a$ expression
• [1] for correct pH = 2.25

(c) [1 mark]

Answer: HCN has the smallest $K_a$ (weakest acid), so its conjugate base ($CN^-$) is the strongest. The weaker the acid, the stronger its conjugate base.

Teaching note: This is an application of the inverse relationship between acid strength and conjugate base strength. $K_a \times K_b = K_w$ for a conjugate pair.

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Question 5 [3 marks]

(a) [2 marks]

Answer: Moles of CH₃COOH initially = $0.0500 \times 0.400 = 0.0200$ mol Moles of NaOH added = $0.0500 \times 0.200 = 0.0100$ mol

The NaOH neutralises half the ethanoic acid: $CH_3COOH + NaOH \rightarrow CH_3COONa + H_2O$

Moles of CH₃COOH remaining = $0.0200 - 0.0100 = 0.0100$ mol Moles of CH₃COO⁻ formed = $0.0100$ mol

The resulting solution contains a mixture of a weak acid (CH₃COOH) and its conjugate base (CH₃COO⁻), which constitutes a buffer solution.

Marking notes:

• [1] for calculating moles and showing partial neutralisation
• [1] for identifying the weak acid/conjugate base pair

(b) [1 mark]

Answer: Since $[CH_3COOH] = [CH_3COO^-]$ (equal moles in the same total volume):

$pH = pK_a + \log\frac{[A^-]}{[HA]} = pK_a + \log(1) = pK_a$ $pK_a = -\log(1.74 \times 10^{-5}) = 4.76$

So pH = 4.76

Teaching note: When the concentrations of weak acid and conjugate base are equal, pH = $pK_a$. This is the half-equivalence point principle.

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Question 6 [3 marks]

(a) [1 mark]

Answer: $Ba(OH)_2(aq) \rightarrow Ba^{2+}(aq) + 2OH^-(aq)$

(b) [2 marks]

Answer: Each formula unit of Ba(OH)₂ produces 2 $OH^-$ ions.

$[OH^-] = 2 \times 0.150 = 0.300 \text{ mol dm}^{-3}$

At 25 °C: $pOH = -\log(0.300) = 0.523$ $pH = 14.00 - 0.523 = 13.48$

Marking notes:

• [1] for $[OH^-] = 0.300$ mol dm⁻³ (recognising the 1:2 stoichiometry)
• [1] for correct pH = 13.48

Common mistakes:

• Forgetting that Ba(OH)₂ releases 2 $OH^-$ ions per formula unit, leading to $[OH^-] = 0.150$ and pH = 13.18.

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Question 7 [2 marks] [2 marks]

Answer: Ammonia is a weak base, so it only partially dissociates in water:

$NH_3(aq) + H_2O(l) \rightleftharpoons NH_4^+(aq) + OH^-(aq)$

• If ammonia were a strong base at 0.100 mol dm⁻³, $[OH^-]$ would be 0.100 mol dm⁻³, giving pH = 13.
• Since dissociation is incomplete, $[OH^-] < 0.100$ mol dm⁻³, so pH < 13.
• The solution is still basic (pH > 7) because some $OH^-$ ions are produced.

Marking notes:

• [1] for explaining partial dissociation and pH < 13
• [1] for explaining pH > 7 (basic solution)

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Question 8 [3 marks]

(a) [1 mark]

Answer: Titrations 1, 2, and 3 are concordant (within 0.10 cm³ of each other). The rough titration is excluded.

$\text{Average titre} = \frac{24.30 + 24.25 + 24.35}{3} = \frac{72.90}{3} = 24.30 \text{ cm}^3$

(b) [1 mark]

Answer: $H_2SO_4 + 2NaOH \rightarrow Na_2SO_4 + 2H_2O$

(c) [1 mark]

Answer: Moles of NaOH = $0.100 \times \frac{24.30}{1000} = 2.43 \times 10^{-3}$ mol

From the equation, mole ratio H₂SO₄ : NaOH = 1 : 2

Moles of H₂SO₄ = $\frac{2.43 \times 10^{-3}}{2} = 1.215 \times 10^{-3}$ mol

$[H_2SO_4] = \frac{1.215 \times 10^{-3}}{0.0250} = 0.0486 \text{ mol dm}^{-3}$

Marking notes:

• Award [1] for correct answer with appropriate working.

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Question 9 [4 marks]

(a) [1 mark]

Answer: $K_{sp} = [Pb^{2+}][I^-]^2$

(b) [1 mark]

Answer: $PbI_2(s) \rightleftharpoons Pb^{2+}(aq) + 2I^-(aq)$

(c) [2 marks]

Answer: Let the solubility of PbI₂ = $s$ mol dm⁻³.

From the equation: $[Pb^{2+}] = s$ and $[I^-] = 2s$

$K_{sp} = s \times (2s)^2 = 4s^3$ $4s^3 = 8.49 \times 10^{-9}$ $s^3 = 2.1225 \times 10^{-9}$ $s = \sqrt[3]{2.1225 \times 10^{-9}} = 1.29 \times 10^{-3} \text{ mol dm}^{-3}$

Marking notes:

• [1] for correct relationship $[I^-] = 2s$ and $K_{sp} = 4s^3$
• [1] for correct answer $s = 1.29 \times 10^{-3}$ mol dm⁻³

Common mistakes:

• Forgetting to square $[I^-]$ in the $K_{sp}$ expression.
• Using $[I^-] = s$ instead of $2s$.

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Question 10 [3 marks]

(a) [1 mark]

Answer: The solution is acidic. NH₄Cl is a salt formed from a weak base (NH₃) and a strong acid (HCl). The $NH_4^+$ ion is the conjugate acid of the weak base and undergoes hydrolysis:

$NH_4^+(aq) + H_2O(l) \rightleftharpoons NH_3(aq) + H_3O^+(aq)$

This produces $H_3O^+$ ions, making the solution acidic.

(b) [2 marks]

Answer: First, find $K_a$ of $NH_4^+$:

$K_a = \frac{K_w}{K_b} = \frac{1.00 \times 10^{-14}}{1.8 \times 10^{-5}} = 5.56 \times 10^{-10} \text{ mol dm}^{-3}$

For the weak acid $NH_4^+$ at 0.020 mol dm⁻³:

$[H^+]^2 = K_a \times [NH_4^+] = 5.56 \times 10^{-10} \times 0.020 = 1.112 \times 10^{-11}$ $[H^+] = \sqrt{1.112 \times 10^{-11}} = 3.33 \times 10^{-6} \text{ mol dm}^{-3}$ $pH = -\log(3.33 \times 10^{-6}) = 5.48$

Marking notes:

• [1] for calculating $K_a$ of $NH_4^+$ using $K_w/K_b$
• [1] for correct pH = 5.48

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Section B

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Question 11 [7 marks]

(a) [2 marks]

Answer: Total volume of solution = $50.0 + 50.0 = 100.0$ cm³

Mass of solution = $100.0 \times 1.00 = 100.0$ g

$q = mc\Delta T = 100.0 \times 4.18 \times 6.8 = 2842.4 \text{ J} = 2.84 \text{ kJ}$

Marking notes:

• [1] for correct mass and substitution
• [1] for correct answer in kJ

(b) [1 mark]

Answer: Moles of HCl = $1.00 \times \frac{50.0}{1000} = 0.0500$ mol Moles of NaOH = $1.00 \times \frac{50.0}{1000} = 0.0500$ mol

From the equation $HCl + NaOH \rightarrow NaCl + H_2O$, the mole ratio is 1:1.

Moles of water formed = 0.0500 mol

(c) [2 marks]

Answer: $\Delta H = -\frac{q}{n} = -\frac{2.84}{0.0500} = -56.8 \text{ kJ mol}^{-1}$

(The negative sign indicates the reaction is exothermic.)

Marking notes:

• [1] for correct calculation
• [1] for correct sign and units

(d) [1 mark]

Answer: Heat loss to the surroundings (or the polystyrene cup is not a perfect insulator / heat absorbed by the cup / incomplete thermal insulation).

(e) [1 mark]

Answer: The temperature rise would be lower. Ethanoic acid is a weak acid and does not fully dissociate. Energy is required to dissociate the ethanoic acid before neutralisation can occur, so less net heat is released. Additionally, the dissociation of the weak acid is endothermic, partially offsetting the exothermic neutralisation.

Marking notes:

• [1] for "lower" with a valid explanation involving weak acid dissociation

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Question 12 [6 marks]

(a) [2 marks]

Answer: From the graph, the pH at the equivalence point is approximately 8.7. Since the pH at the equivalence point is greater than 7, this confirms that HA is a weak acid. The salt formed (NaA) is the salt of a weak acid and strong base, so the conjugate base $A^-$ hydrolyses to produce an alkaline solution.

Marking notes:

• [1] for reading pH ≈ 8.7 from the graph
• [1] for concluding HA is a weak acid with explanation

(b) [2 marks]

Answer: At the half-equivalence point, half the acid has been neutralised, so $[HA] = [A^-]$ and $pH = pK_a$.

From the graph, the half-equivalence point occurs at 12.5 cm³ NaOH (half of 25.0 cm³), where the pH ≈ 4.8.

$pK_a = 4.8$ $K_a = 10^{-4.8} = 1.58 \times 10^{-5} \text{ mol dm}^{-3}$

Marking notes:

• [1] for identifying the half-equivalence point and that pH = $pK_a$
• [1] for correct $K_a$ value

(c) [2 marks]

Answer: Phenolphthalein changes colour in the pH range 8.2–10.0 (colourless to pink). From the graph, the steep vertical portion of the titration curve (the pH jump) occurs between approximately pH 7 and pH 10, which falls within the phenolphthalein transition range. Therefore, the colour change will occur sharply at the equivalence point, making phenolphthalein a suitable indicator.

Marking notes:

• [1] for stating the pH range of phenolphthalein
• [1] for linking the indicator range to the steep portion of the curve

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Question 13 [6 marks]

(a) [3 marks]

Answer: Using the Henderson–Hasselbalch equation:

$pH = pK_a + \log\frac{[CH_3COO^-]}{[CH_3COOH]}$

$pK_a = -\log(1.74 \times 10^{-5}) = 4.76$

$5.00 = 4.76 + \log\frac{[CH_3COO^-]}{[CH_3COOH]}$

$\log\frac{[CH_3COO^-]}{[CH_3COOH]} = 5.00 - 4.76 = 0.24$

$\frac{[CH_3COO^-]}{[CH_3COOH]} = 10^{0.24} = 1.74$

Marking notes:

• [1] for correct $pK_a$ calculation
• [1] for correct substitution into Henderson–Hasselbalch equation
• [1] for correct ratio = 1.74

(b) [3 marks]

Answer: Moles of CH₃COOH = $0.100 \times 0.500 = 0.0500$ mol

Since $\frac{[CH_3COO^-]}{[CH_3COOH]} = 1.74$ and both are in the same volume:

$\frac{n_{CH_3COO^-}}{n_{CH_3COOH}} = 1.74$

$n_{CH_3COO^-} = 1.74 \times 0.0500 = 0.0870 \text{ mol}$

Mass of CH₃COONa = $0.0870 \times 82.0 = 7.13$ g

Marking notes:

• [1] for moles of CH₃COOH = 0.0500 mol
• [1] for moles of CH₃COONa = 0.0870 mol
• [1] for correct mass = 7.13 g

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Question 14 [5 marks]

(a) [3 marks]

Answer: Moles of NaOH = $0.0300 \times 0.200 = 6.00 \times 10^{-3}$ mol Moles of HCl = $0.0200 \times 0.300 = 6.00 \times 10^{-3}$ mol

From the equation: $NaOH + HCl \rightarrow NaCl + H_2O$ (1:1 ratio)

Since moles of NaOH = moles of HCl, the reaction is exactly at the equivalence point. Neither reagent is in excess.

Wait — let me recheck: Moles of NaOH = $0.0300 \times 0.200 = 0.00600$ mol. Moles of HCl = $0.0200 \times 0.300 = 0.00600$ mol. These are equal, so the solution is neutral.

Correction for the question design: Let me recalculate with the given values.

Actually, the values as given produce exact equivalence. For the purpose of this question, the answer is:

Moles of NaOH = $0.0300 \times 0.200 = 6.00 \times 10^{-3}$ mol Moles of HCl = $0.0200 \times 0.300 = 6.00 \times 10^{-3}$ mol

The reaction is: $NaOH + HCl \rightarrow NaCl + H_2O$

Since the mole ratio is 1:1 and moles are equal, neither reagent is in excess. The resulting solution contains only NaCl (a neutral salt) and water.

Marking notes:

• [1] for correct moles of NaOH
• [1] for correct moles of HCl
• [1] for correct conclusion that neither is in excess

(b) [2 marks]

Answer: Since neither reagent is in excess and the salt NaCl is formed from a strong acid (HCl) and strong base (NaOH), the solution is neutral.

$pH = 7.00$

Marking notes:

• [1] for stating the solution is neutral
• [1] for pH = 7.00

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Question 15 [6 marks]

(a) [2 marks]

Answer: $Ca(OH)_2(s) \rightleftharpoons Ca^{2+}(aq) + 2OH^-(aq)$

$K_{sp} = [Ca^{2+}][OH^-]^2$

Marking notes:

• [1] for correct equilibrium equation
• [1] for correct $K_{sp}$ expression

(b) [2 marks]

Answer: Adding NaOH increases the concentration of $OH^-$ ions in solution. According to Le Chatelier's principle, the equilibrium will shift to the left (towards the solid Ca(OH)₂) to counteract the increase in $[OH^-]$. This causes more Ca(OH)₂ to precipitate, decreasing its solubility.

Marking notes:

• [1] for identifying the shift to the left
• [1] for stating that solubility decreases

(c) [2 marks]

Answer: $pH = 12.35$

$pOH = 14.00 - 12.35 = 1.65$ $[OH^-] = 10^{-1.65} = 2.24 \times 10^{-2} \text{ mol dm}^{-3}$

From the dissolution equation: $[Ca^{2+}] = \frac{1}{2}[OH^-] = \frac{2.24 \times 10^{-2}}{2} = 1.12 \times 10^{-2}$ mol dm⁻³

$K_{sp} = [Ca^{2+}][OH^-]^2 = (1.12 \times 10^{-2})(2.24 \times 10^{-2})^2$ $K_{sp} = 1.12 \times 10^{-2} \times 5.02 \times 10^{-4} = 5.62 \times 10^{-6} \text{ mol}^3 \text{ dm}^{-9}$

Marking notes:

• [1] for correct $[OH^-]$ and $[Ca^{2+}]$ values
• [1] for correct $K_{sp} = 5.62 \times 10^{-6}$ mol³ dm⁻⁹

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End of Answer Key

Section A Total: 30 marks | Section B Total: 30 marks | Grand Total: 60 marks