A Level H2 Chemistry Practice Paper 3

TuitionGoWhere Practice Paper — Chemistry H2 A-Level

Answer Key — Version 3 of 5

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Section A: Multiple Choice and Short Answer (Questions 1–10)

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Question 1 [1]

Answer: B

$K_w = [H^+][OH^-]$

Explanation: The ionic product of water is defined as the product of the concentrations of hydrogen ions and hydroxide ions in aqueous solution. At 25 °C, $K_w = 1.00 \times 10^{-14} \text{ mol}^2 \text{ dm}^{-6}$. This is a fundamental definition that students must recall. Option A is incorrect because water is a pure liquid and its concentration is not included in the equilibrium expression. Option C is a ratio, not a product. Option D is a sum, not a product.

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Question 2 [1]

Answer: A — $2.51 \times 10^{-11} \text{ mol dm}^{-3}$

Working:

Given: pH = 3.40

Step 1: Calculate $[H^+]$ $[H^+] = 10^{-\text{pH}} = 10^{-3.40} = 3.98 \times 10^{-4} \text{ mol dm}^{-3}$

Step 2: Use $K_w$ to find $[OH^-]$ $K_w = [H^+][OH^-] = 1.00 \times 10^{-14}$ $[OH^-] = \frac{1.00 \times 10^{-14}}{3.98 \times 10^{-4}} = 2.51 \times 10^{-11} \text{ mol dm}^{-3}$

Explanation: This question tests the relationship between pH, $[H^+]$, and $[OH^-]$ via $K_w$. Students must first convert pH to $[H^+]$ using the antilog, then use the $K_w$ expression to find $[OH^-]$. A common mistake is to calculate $10^{-3.40}$ incorrectly or to confuse $[H^+]$ with $[OH^-]$.

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Question 3 [2]

(a) [1] A Brønsted–Lowry acid is a proton ($H^+$) donor.

Explanation: The Brønsted–Lowry theory defines acids and bases in terms of proton transfer. An acid donates a proton; a base accepts a proton. This is the most commonly used definition in A-Level Chemistry.

(b) [1] A buffer solution is a solution that resists changes in pH when small amounts of acid or base are added (or when it is diluted).

Explanation: A buffer typically consists of a weak acid and its conjugate base (or a weak base and its conjugate acid). The key property is resistance to pH change, not maintenance of a constant pH.

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Question 4 [2]

Answer: pH = 12.40

Working:

$NaOH$ is a strong base and dissociates completely: $NaOH \rightarrow Na^+ + OH^-$

$[OH^-] = 0.025 \text{ mol dm}^{-3}$

Step 1: Calculate pOH $\text{pOH} = -\log_{10}[OH^-] = -\log_{10}(0.025) = 1.60$

Step 2: Calculate pH $\text{pH} = 14.00 - \text{pOH} = 14.00 - 1.60 = 12.40$

Alternatively: $[H^+] = \frac{K_w}{[OH^-]} = \frac{1.00 \times 10^{-14}}{0.025} = 4.00 \times 10^{-13} \text{ mol dm}^{-3}$ $\text{pH} = -\log_{10}(4.00 \times 10^{-13}) = 12.40$

Explanation: For strong base pH calculations, students can either calculate pOH first and subtract from 14, or calculate $[H^+]$ from $K_w$ and then find pH. Both methods are valid. The key is recognising that $NaOH$ dissociates completely, so $[OH^-] = [NaOH]$.

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Question 5 [3]

(a) [1] $K_a = \frac{[CH_3COO^-][H^+]}{[CH_3COOH]}$

Explanation: The $K_a$ expression is written as the product of the concentrations of the products (dissociated ions) divided by the concentration of the undissociated acid. Water is omitted as it is a pure liquid.

(b) [2] Answer: pH = 2.88

Working:

Assumption: The dissociation of ethanoic acid is small, so $[CH_3COOH] \approx 0.100 \text{ mol dm}^{-3}$ at equilibrium.

$K_a = \frac{[CH_3COO^-][H^+]}{[CH_3COOH]} = \frac{[H^+]^2}{0.100} = 1.74 \times 10^{-5}$

Since $[CH_3COO^-] = [H^+]$ (1:1 stoichiometry of dissociation):

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

Marking notes:

• 1 mark for correct $K_a$ expression or correct method setup
• 1 mark for correct pH value

Explanation: For weak acid pH calculations, the key assumption is that the degree of dissociation is small, so the equilibrium concentration of the undissociated acid is approximately equal to the initial concentration. This allows the simplification $[H^+]^2 = K_a \times C$. Students should always state this assumption. The assumption is valid if the percentage dissociation is less than 5%.

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Question 6 [3]

(a) [2] Answer: pH = 4.76

Working:

After mixing equal volumes, the concentrations are halved: $[CH_3COOH] = \frac{0.200}{2} = 0.100 \text{ mol dm}^{-3}$ $[CH_3COO^-] = \frac{0.200}{2} = 0.100 \text{ mol dm}^{-3}$

Using the Henderson–Hasselbalch equation: $\text{pH} = \text{p}K_a + \log_{10}\frac{[\text{salt}]}{[\text{acid}]}$ $\text{p}K_a = -\log_{10}(1.74 \times 10^{-5}) = 4.76$ $\text{pH} = 4.76 + \log_{10}\frac{0.100}{0.100} = 4.76 + \log_{10}(1) = 4.76 + 0 = 4.76$

Marking notes:

• 1 mark for correct pKa calculation or recognition that [salt]/[acid] = 1
• 1 mark for correct final pH

(b) [1] When a small amount of $HCl$ is added, the $H^+$ ions react with the ethanoate ions ($CH_3COO^-$) to form more ethanoic acid ($CH_3COOH$). The ratio $[CH_3COO^-]/[CH_3COOH]$ changes only slightly, so the pH remains approximately constant.

Explanation: The buffer works because the added $H^+$ is consumed by the conjugate base ($CH_3COO^-$) present in the buffer. This is the fundamental mechanism of buffer action. The pH change is small because the logarithmic relationship means that even significant changes in the ratio produce only small pH changes.

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Question 7 [3]

(a) [1] From the graph, the pH at the equivalence point (25.0 cm³) is approximately 8.8 (accept 8.5–9.0).

Explanation: The equivalence point is where the steepest part of the curve occurs, at 25.0 cm³ of NaOH added. Students should read the pH value at this point from the y-axis.

(b) [1] At the equivalence point, all the ethanoic acid has been converted to sodium ethanoate ($CH_3COONa$). The ethanoate ion is the conjugate base of a weak acid and undergoes hydrolysis in water: $CH_3COO^- + H_2O \rightleftharpoons CH_3COOH + OH^-$ This produces $OH^-$ ions, making the solution alkaline (pH > 7).

Explanation: The pH at the equivalence point of a weak acid–strong base titration is always greater than 7 because the salt formed contains the conjugate base of the weak acid, which hydrolyses to produce hydroxide ions.

(c) [1] Phenolphthalein. The equivalence point pH (~8.8) falls within the pH range of phenolphthalein (8.3–10.0), so the colour change (colourless to pink) will occur at the equivalence point.

Explanation: The indicator must have a transition range that includes the equivalence point pH. For weak acid–strong base titrations, the equivalence point is in the alkaline range, so phenolphthalein is suitable. Methyl orange would change colour too early (in the buffer region) and is not suitable.

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Question 8 [3]

(a) [1] Titration 1 (24.40 cm³) and Titration 3 (24.30 cm³) are concordant (within 0.10 cm³ of each other). Titration 2 (48.70 cm³) appears to be a cumulative reading (the student did not reset the burette to 0.00 between titrations 1 and 2). The rough titration (24.80 cm³) is not used in the mean calculation. The concordant results are titrations 1 and 3 (or 2 and 3 if titration 2 is read as 24.30 cm³ from the difference 48.70 − 24.40 = 24.30).

Re-evaluation: Looking more carefully:

• Titration 2: Final = 48.70, Initial = 24.40 → Volume = 48.70 − 24.40 = 24.30 cm³
• Titration 3: Final = 24.30, Initial = 0.00 → Volume = 24.30 cm³

So titrations 2 and 3 are concordant at 24.30 cm³. Titration 1 (24.40 cm³) is also concordant (within 0.10 cm³).

Answer: All three accurate titrations (1, 2, and 3) are concordant. No anomalous result among the accurate titrations. The rough titration is not included in the mean.

(b) [1] Mean titre $= \frac{24.40 + 24.30 + 24.30}{3} = \frac{73.00}{3} = 24.33 \text{ cm}^3$

(c) [1] $HCl + NaOH \rightarrow NaCl + H_2O$

Moles of $NaOH = 0.100 \times \frac{24.33}{1000} = 2.433 \times 10^{-3} \text{ mol}$

From the 1:1 stoichiometry, moles of $HCl = 2.433 \times 10^{-3} \text{ mol}$

$[HCl] = \frac{2.433 \times 10^{-3}}{\frac{25.0}{1000}} = \frac{2.433 \times 10^{-3}}{0.0250} = 0.0973 \text{ mol dm}^{-3}$

Explanation: This is a standard acid–titration calculation. Students must identify concordant titres (within 0.10 cm³), calculate the mean, then use the stoichiometry of the reaction to find the unknown concentration.

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Question 9 [3]

(a) [1] $K_{sp} = [Mg^{2+}][OH^-]^2$

Explanation: The $K_{sp}$ expression is written from the dissolution equilibrium: $Mg(OH)_2(s) \rightleftharpoons Mg^{2+}(aq) + 2OH^-(aq)$ The concentration of the solid is omitted, and the stoichiometric coefficients become exponents.

(b) [2] Answer: $1.65 \times 10^{-4} \text{ mol dm}^{-3}$

Working:

Let the solubility of $Mg(OH)_2 = s \text{ mol dm}^{-3}$

Then: $[Mg^{2+}] = s$ and $[OH^-] = 2s$

$K_{sp} = [Mg^{2+}][OH^-]^2 = s \times (2s)^2 = 4s^3$ $4s^3 = 1.8 \times 10^{-11}$ $s^3 = \frac{1.8 \times 10^{-11}}{4} = 4.5 \times 10^{-12}$ $s = \sqrt[3]{4.5 \times 10^{-12}} = 1.65 \times 10^{-4} \text{ mol dm}^{-3}$

Marking notes:

• 1 mark for correct setup ($K_{sp} = 4s^3$)
• 1 mark for correct final answer

Explanation: For solubility product calculations, students must relate the molar solubility $s$ to the ion concentrations using the stoichiometry of the dissolution equation. A common mistake is to forget that $[OH^-] = 2s$ (not $s$), which leads to $K_{sp} = 4s^3$ rather than $K_{sp} = s^3$.

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Question 10 [4]

(a) [2] Hydrochloric acid is a strong acid and dissociates completely, contributing $[H^+] = 0.100 \text{ mol dm}^{-3}$. Ethanoic acid is a weak acid with $K_a = 1.74 \times 10^{-5}$, so its contribution to $[H^+]$ is very small compared to $HCl$. Additionally, the high $[H^+]$ from $HCl suppresses the dissociation of ethanoic acid (common ion effect). Therefore, the total$[H^+] \approx 0.100 \text{ mol dm}^{-3}$, giving$\text{pH} = -\log_{10}(0.100) = 1.0$.

Explanation: In a mixture of a strong acid and a weak acid, the pH is dominated by the strong acid because it dissociates completely. The common ion effect ($H^+$ from $HCl$) further suppresses the dissociation of the weak acid.

(b) [2] Answer: $[CH_3COO^-] = 1.74 \times 10^{-5} \text{ mol dm}^{-3}$

Working:

For ethanoic acid: $K_a = \frac{[CH_3COO^-][H^+]}{[CH_3COOH]} = 1.74 \times 10^{-5}$

In the mixture: $[H^+] \approx 0.100$ (from $HCl$) and $[CH_3COOH] \approx 0.100$ (undissociated)

$1.74 \times 10^{-5} = \frac{[CH_3COO^-] \times 0.100}{0.100}$ $[CH_3COO^-] = 1.74 \times 10^{-5} \text{ mol dm}^{-3}$

Marking notes:

• 1 mark for correct substitution into $K_a$ expression
• 1 mark for correct final answer

Explanation: This question demonstrates the common ion effect quantitatively. The presence of $H^+$ from the strong acid suppresses the dissociation of the weak acid to such an extent that $[CH_3COO^-]$ is extremely small — equal to $K_a$ itself when $[H^+] = [CH_3COOH]$.

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Section B: Structured and Data-Based Questions (Questions 11–20)

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Question 11 [4]

(a) [1] Increasing acid strength: $HCOOH < HNO_2 < HF$

(Acids with larger $K_a$ values are stronger acids.)

Explanation: The acid dissociation constant $K_a$ is a measure of acid strength. A larger $K_a$ means greater dissociation and therefore a stronger acid. $HF$ has the largest $K_a$ ($6.8 \times 10^{-4}$) and is the strongest; $HCOOH$ has the smallest $K_a$ ($1.8 \times 10^{-4}$) and is the weakest.

(b) [2] Answer: pH = 2.52

Working:

$K_a = \frac{[H^+]^2}{[HCOOH]} = \frac{[H^+]^2}{0.050} = 1.8 \times 10^{-4}$ $[H^+]^2 = 1.8 \times 10^{-4} \times 0.050 = 9.0 \times 10^{-6}$ $[H^+] = \sqrt{9.0 \times 10^{-6}} = 3.0 \times 10^{-3} \text{ mol dm}^{-3}$ $\text{pH} = -\log_{10}(3.0 \times 10^{-3}) = 2.52$

Marking notes:

• 1 mark for correct method/equation setup
• 1 mark for correct pH value

(c) [1] Methanoic acid ($HCOOH$) would produce the highest pH. It is the weakest acid (smallest $K_a$), so it dissociates the least, producing the lowest $[H^+]$ and therefore the highest pH.

Explanation: For solutions of equal concentration, the weakest acid produces the fewest $H^+$ ions and therefore has the highest pH. This is the inverse relationship between acid strength and pH.

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Question 12 [5]

(a) [2] Answer: $[CH_3COONa] = 0.200 \text{ mol dm}^{-3}$

Working:

Moles of $CH_3COONa = \frac{4.10}{82.0} = 0.0500 \text{ mol}$

$[CH_3COONa] = \frac{0.0500}{\frac{250}{1000}} = \frac{0.0500}{0.250} = 0.200 \text{ mol dm}^{-3}$

Marking notes:

• 1 mark for correct moles calculation
• 1 mark for correct concentration

(b) [2] Answer: pH = 5.06

Working:

$\text{p}K_a = -\log_{10}(1.74 \times 10^{-5}) = 4.76$ $\text{pH} = \text{p}K_a + \log_{10}\frac{[\text{salt}]}{[\text{acid}]} = 4.76 + \log_{10}\frac{0.200}{0.100} = 4.76 + \log_{10}(2) = 4.76 + 0.30 = 5.06$

Marking notes:

• 1 mark for correct pKa and log ratio calculation
• 1 mark for correct final pH

(c) [1] The pH change will be small. The buffer contains sufficient $CH_3COO^-$ to react with the added $H^+$ ions, converting them to $CH_3COOH$. The ratio $[CH_3COO^-]/[CH_3COOH]$ changes only slightly, so the pH remains relatively constant.

Explanation: This tests understanding of buffer action. The added acid is neutralised by the conjugate base component of the buffer, minimising the pH change.

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Question 13 [4]

(a) [2] Ammonium chloride is a salt of a weak base ($NH_3$) and a strong acid ($HCl$). The ammonium ion ($NH_4^+$) is the conjugate acid of the weak base and undergoes hydrolysis in water: $NH_4^+ + H_2O \rightleftharpoons NH_3 + H_3O^+$ This produces $H_3O^+$ (or $H^+$) ions, making the solution acidic (pH < 7).

Marking notes:

• 1 mark for identifying the salt as from weak base + strong acid
• 1 mark for correct hydrolysis equation and explanation

(b) [2] Sodium ethanoate is a salt of a strong base ($NaOH$) and a weak acid ($CH_3COOH$). The ethanoate ion ($CH_3COO^-$) is the conjugate base of the weak acid and undergoes hydrolysis in water: $CH_3COO^- + H_2O \rightleftharpoons CH_3COOH + OH^-$ This produces $OH^-$ ions, making the solution alkaline (pH > 7).

Marking notes:

• 1 mark for identifying the salt as from strong base + weak acid
• 1 mark for correct hydrolysis equation and explanation

Explanation: The acidity or alkalinity of a salt solution depends on the relative strengths of the parent acid and base. Salts of strong acid + weak base are acidic; salts of weak acid + strong base are alkaline; salts of strong acid + strong base are neutral.

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Question 14 [4]

(a) [2] The mixture contains two acids: $HCl$ (a strong monoprotic acid) and $H_2SO_4$ (a strong diprotic acid). The first equivalence point corresponds to the neutralisation of the first proton from $H_2SO_4$ plus all the $H^+$ from $HCl$. The second equivalence point corresponds to the neutralisation of the second proton from $H_2SO_4$. Alternatively, the two equivalence points arise because $H_2SO_4$ can donate two protons sequentially.

More precisely: The first equivalence point at ~15.0 cm³ corresponds to neutralisation of all $HCl$ and the first proton of $H_2SO_4$. The second equivalence point at ~30.0 cm³ corresponds to neutralisation of the second proton of $H_2SO_4$. The volume between the two equivalence points (15.0 cm³) equals the volume to the first equivalence point, confirming that the second proton of $H_2SO_4$ requires the same amount of base as the first.

Marking notes:

• 1 mark for identifying the two acids and their nature
• 1 mark for explaining the two-stage neutralisation

(b) [1] From the graph, the second equivalence point occurs at 30.0 cm³ of $NaOH$.

(c) [2] Answer: $[H^+]_{\text{total}} = 0.120 \text{ mol dm}^{-3}$

Working:

Total moles of $NaOH$ used to second equivalence point: $n_{NaOH} = 0.100 \times \frac{30.0}{1000} = 3.00 \times 10^{-3} \text{ mol}$

Since each mole of $NaOH$ neutralises one mole of $H^+$: $\text{Total moles of } H^+ = 3.00 \times 10^{-3} \text{ mol}$

$[H^+]_{\text{total}} = \frac{3.00 \times 10^{-3}}{\frac{25.0}{1000}} = \frac{3.00 \times 10^{-3}}{0.0250} = 0.120 \text{ mol dm}^{-3}$

Marking notes:

• 1 mark for correct moles of NaOH calculation
• 1 mark for correct total H+ concentration

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Question 15 [4]

(a) [1] The common ion effect is the reduction in the solubility of a sparingly soluble salt when a soluble compound containing a common ion is added to the solution. It is an application of Le Chatelier's principle.

Explanation: The common ion effect occurs because the addition of a common ion shifts the dissolution equilibrium to the left (towards the solid), reducing solubility.

(b)(i) [1] Answer: $1.34 \times 10^{-5} \text{ mol dm}^{-3}$

Working:

$AgCl(s) \rightleftharpoons Ag^+(aq) + Cl^-(aq)$ $K_{sp} = [Ag^+][Cl^-] = s^2 = 1.8 \times 10^{-10}$ $s = \sqrt{1.8 \times 10^{-10}} = 1.34 \times 10^{-5} \text{ mol dm}^{-3}$

(b)(ii) [2] Answer: $1.8 \times 10^{-9} \text{ mol dm}^{-3}$

Working:

In $0.100 \text{ mol dm}^{-3}$ $NaCl$, $[Cl^-] = 0.100 \text{ mol dm}^{-3}$ (from the soluble salt).

Let the solubility of $AgCl$ in this solution = $s'$

Then: $[Ag^+] = s'$ and $[Cl^-] = 0.100 + s' \approx 0.100$ (since $s' \ll 0.100$)

$K_{sp} = [Ag^+][Cl^-] = s' \times 0.100 = 1.8 \times 10^{-10}$ $s' = \frac{1.8 \times 10^{-10}}{0.100} = 1.8 \times 10^{-9} \text{ mol dm}^{-3}$

Marking notes:

• 1 mark for correct setup with common ion
• 1 mark for correct final answer

Explanation: The solubility of $AgCl$ is dramatically reduced (by a factor of ~7400) in the presence of the common ion $Cl^-$. This is a clear demonstration of the common ion effect. The approximation $0.100 + s' \approx 0.100$ is valid because $s'$ is negligible compared to 0.100.

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Question 16 [4]

(a) [2] As temperature increases, $K_w$ increases. This means the dissociation of water: $H_2O(l) \rightleftharpoons H^+(aq) + OH^-(aq)$ is favoured at higher temperatures. By Le Chatelier's principle, if the forward reaction is favoured by increasing temperature, the reaction must be endothermic ($\Delta H > 0$). The dissociation of water absorbs heat, so increasing temperature shifts the equilibrium to the right, producing more $H^+$ and $OH^-$ ions and increasing $K_w$.

Marking notes:

• 1 mark for stating the reaction is endothermic
• 1 mark for correct application of Le Chatelier's principle

(b) [1] Answer: pH = 6.51

Working:

At 60 °C: $K_w = 9.61 \times 10^{-14}$

In pure water: $[H^+] = [OH^-] = \sqrt{K_w} = \sqrt{9.61 \times 10^{-14}} = 3.10 \times 10^{-7} \text{ mol dm}^{-3}$

$\text{pH} = -\log_{10}(3.10 \times 10^{-7}) = 6.51$

(c) [1] Pure water at 60 °C is neutral. In pure water, $[H^+] = [OH^-]$ at all temperatures. Although the pH is less than 7, this does not mean the solution is acidic — it simply reflects the higher $K_w$ at elevated temperature. Neutrality is defined by $[H^+] = [OH^-]$, not by pH = 7.

Explanation: This is a common misconception. pH = 7 is only neutral at 25 °C. At higher temperatures, the neutral pH is lower because $K_w$ is larger. Students must understand that neutrality means $[H^+] = [OH^-]$, not pH = 7.

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Question 17 [3]

(a) [2] Answer: $K_a = 2.0 \times 10^{-5} \text{ mol dm}^{-3}$

Working:

$[H^+] = 10^{-\text{pH}} = 10^{-3.00} = 1.0 \times 10^{-3} \text{ mol dm}^{-3}$

For the dissociation: $HA \rightleftharpoons H^+ + A^-$

At equilibrium: $[H^+] = [A^-] = 1.0 \times 10^{-3}$ and $[HA] = 0.050 - 1.0 \times 10^{-3} \approx 0.049$

$K_a = \frac{[H^+][A^-]}{[HA]} = \frac{(1.0 \times 10^{-3})^2}{0.049} = \frac{1.0 \times 10^{-6}}{0.049} = 2.04 \times 10^{-5} \approx 2.0 \times 10^{-5} \text{ mol dm}^{-3}$

Marking notes:

• 1 mark for correct $[H^+]$ calculation and equilibrium concentrations
• 1 mark for correct $K_a$ value

(b) [1] Answer: 2.0%

Working:

$\text{Percentage dissociation} = \frac{[H^+]_{\text{equilibrium}}}{[HA]_{\text{initial}}} \times 100\% = \frac{1.0 \times 10^{-3}}{0.050} \times 100\% = 2.0\%$

Explanation: Percentage dissociation indicates what fraction of the acid molecules have dissociated. For weak acids, this is typically small (< 5%), which validates the approximation used in many $K_a$ calculations.

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Question 18 [4]

(a) [3] To prepare a buffer with pH = 5.00 using ethanoic acid and sodium ethanoate:

Step 1: Calculate the required ratio using the Henderson–Hasselbalch equation. $\text{pH} = \text{p}K_a + \log_{10}\frac{[CH_3COO^-]}{[CH_3COOH]}$ $5.00 = 4.76 + \log_{10}\frac{[CH_3COO^-]}{[CH_3COOH]}$ $\log_{10}\frac{[CH_3COO^-]}{[CH_3COOH]} = 0.24$ $\frac{[CH_3COO^-]}{[CH_3COOH]} = 10^{0.24} = 1.74$

Step 2: The ratio of $[CH_3COO^-]$ to $[CH_3COOH]$ should be 1.74 : 1.

Step 3: To prepare the buffer, mix appropriate volumes/concentrations of ethanoic acid and sodium ethanoate solutions such that this ratio is achieved. For example, mix $100 \text{ cm}^3$ of $0.100 \text{ mol dm}^{-3}$ ethanoic acid with $174 \text{ cm}^3$ of $0.100 \text{ mol dm}^{-3}$ sodium ethanoate (or any combination giving the 1.74:1 ratio). Verify the pH with a pH meter and adjust if necessary.

Marking notes:

• 1 mark for correct pKa calculation
• 1 mark for correct ratio calculation
• 1 mark for practical preparation method

(b) [1] A buffer resists pH changes because it contains significant amounts of both a weak acid and its conjugate base. When acid is added, the conjugate base neutralises it; when base is added, the weak acid neutralises it. The ratio of acid to conjugate base changes only slightly, so the pH remains relatively constant.

Explanation: The buffer action relies on the equilibrium between the weak acid and its conjugate base. Both components must be present in sufficient quantities to neutralise added acid or base.

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Question 19 [3]

(a) [1] $CO_3^{2-} + H_2O \rightleftharpoons HCO_3^- + OH^-$

Explanation: The carbonate ion is the conjugate base of the weak acid $HCO_3^-$. It accepts a proton from water, producing hydroxide ions.

(b) [1] The hydrolysis of $CO_3^{2-}$ produces $OH^-$ ions, which increases the $[OH^-]$ in the solution, making it alkaline (pH > 7).

Explanation: Any reaction that produces $OH^-$ ions will make the solution alkaline. The carbonate ion is a relatively strong conjugate base (since $HCO_3^-$ is a weak acid), so the hydrolysis is significant.

(c) [1] A solution of $NaHCO_3$ would be alkaline (pH > 7). The hydrogencarbonate ion ($HCO_3^-$) can act as both an acid (donating $H^+$ to form $CO_3^{2-}$) and a base (accepting $H^+$ to form $H_2CO_3$). However, the $K_b$ of $HCO_3^-$ (as a base, $K_b = K_w/K_{a1}$) is greater than the $K_a$ of $HCO_3^-$ (as an acid, $K_{a2}$), so the solution is slightly alkaline.

Explanation: $HCO_3^-$ is amphiprotic. Whether its solution is acidic or alkaline depends on whether its acidic or basic character dominates. Since $K_b(HCO_3^-) > K_a(HCO_3^-)$, the solution is alkaline (pH ≈ 8.3 for typical concentrations).

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Question 20 [5]

(a) [2]

The concordant titres are titrations 1, 2, and 3 (all within 0.10 cm³ of each other: 25.85, 25.75, 25.80).

Mean titre $= \frac{25.85 + 25.75 + 25.80}{3} = \frac{77.40}{3} = 25.80 \text{ cm}^3$

Marking notes:

• 1 mark for identifying concordant titres
• 1 mark for correct mean calculation

(b) [2] Answer: $[H_2SO_4] = 0.0645 \text{ mol dm}^{-3}$

Working:

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

Moles of $NaOH = 0.100 \times \frac{25.80}{1000} = 2.58 \times 10^{-3} \text{ mol}$

From the 1:2 stoichiometry: $\text{Moles of } H_2SO_4 = \frac{2.58 \times 10^{-3}}{2} = 1.29 \times 10^{-3} \text{ mol}$

$[H_2SO_4] = \frac{1.29 \times 10^{-3}}{\frac{20.0}{1000}} = \frac{1.29 \times 10^{-3}}{0.0200} = 0.0645 \text{ mol dm}^{-3}$

Marking notes:

• 1 mark for correct stoichiometry (1:2 ratio)
• 1 mark for correct final concentration

(c) [1] Methyl orange is suitable because the titration of a strong acid ($H_2SO_4$) with a strong base ($NaOH$) has an equivalence point at pH 7. Methyl orange changes colour in the pH range 3.1–4.4, which is on the acidic side of the steep portion of the titration curve. The colour change (yellow to orange/red) occurs sharply at the equivalence point for this strong acid–strong base titration.

Note: In practice, either methyl orange or phenolphthalein can be used for strong acid–strong base titrations because the pH change at the equivalence point is very steep (pH 3–11), so both indicators change colour within this range.

Explanation: For strong acid–strong base titrations, the equivalence point is at pH 7 and the pH change is very steep. Both methyl orange and phenolphthalein are suitable because their transition ranges fall within the steep portion of the titration curve.

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Summary of Marks

Section	Questions	Marks
A	1–10	25
B	11–20	25
Total		50