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A Level H1 Chemistry Practice Paper 3

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TuitionGoWhere Exam Practice (AI)

A-Level Chemistry H1 Practice Paper - Version 3

Topic: Acids, Bases and Salts

Subject: Chemistry
Level: A-Level H1
Paper: Practice Paper (Topic Focus)
Duration: 1 hour
Total Marks: 40

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 in this question paper.
The number of marks is given in brackets [ ] at the end of each question or part question.
You may use a scientific calculator.
A Data Booklet is provided for reference.

Section A: Structured Questions

1. Ethanoic acid, CH₃COOH, is a weak acid commonly found in vinegar.

(a) Define the term weak acid.
[1]
........................................................................................................................................
........................................................................................................................................

(b) Write an equation, including state symbols, to show the dissociation of ethanoic acid in water.
[1]
........................................................................................................................................

(c) Explain, in terms of bonding and structure, why ethanoic acid has a higher boiling point than ethanol, despite having a similar molar mass.
[2]
........................................................................................................................................
........................................................................................................................................
........................................................................................................................................

2. A student titrated 25.0 cm³ of a solution of benzoic acid, C₆H₅COOH, with 0.100 mol dm⁻³ sodium hydroxide, NaOH. The pH of the solution was monitored throughout the titration.

(a) The initial pH of the benzoic acid solution was 2.90. Calculate the concentration of H⁺ ions in this solution.
[1]
........................................................................................................................................

(b) The equivalence point was reached when 20.0 cm³ of NaOH had been added. Calculate the initial concentration of the benzoic acid solution.
[2]
........................................................................................................................................
........................................................................................................................................
........................................................................................................................................

pH
14 |
   |
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 0 |________________________________________ Volume of NaOH / cm³
    0                                    40

3. Buffer solutions are essential in maintaining stable pH conditions in biological systems.

(a) State what is meant by a buffer solution.
[1]
........................................................................................................................................
........................................................................................................................................

(b) A buffer solution is prepared by mixing 0.10 mol dm⁻³ ethanoic acid with 0.10 mol dm⁻³ sodium ethanoate.
Given that $K_a$ for ethanoic acid is $1.7 \times 10^{-5}$ mol dm⁻³, calculate the pH of this buffer solution.
[2]
........................................................................................................................................
........................................................................................................................................
........................................................................................................................................

(c) Explain how this buffer solution resists a change in pH when a small amount of strong acid (H⁺) is added. Include an equation in your answer.
[2]
........................................................................................................................................
........................................................................................................................................
........................................................................................................................................

4. Aluminium oxide, Al₂O₃, is described as an amphoteric oxide.

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

(b) Write balanced ionic equations for the reaction of aluminium oxide with:
(i) Dilute hydrochloric acid.
[1]
........................................................................................................................................

(ii) Aqueous sodium hydroxide.
[1]
........................................................................................................................................

5. The solubility product constant, $K_{sp}$ , is used to describe the equilibrium between a solid salt and its ions in a saturated solution.

(a) Write the expression for the solubility product, $K_{sp}$ , for magnesium hydroxide, Mg(OH)₂.
[1]
........................................................................................................................................

(b) The value of $K_{sp}$ for Mg(OH)₂ is $2.0 \times 10^{-11}$ mol³ dm⁻⁹ at 298 K. Calculate the solubility of Mg(OH)₂ in mol dm⁻³.
[2]
........................................................................................................................................
........................................................................................................................................
........................................................................................................................................

(c) Explain why the solubility of Mg(OH)₂ decreases when it is placed in a solution of magnesium chloride, MgCl₂.
[2]
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........................................................................................................................................
........................................................................................................................................

Section B: Data Interpretation and Application

6. Rainwater is naturally slightly acidic due to the presence of dissolved carbon dioxide. In industrial areas, rainwater can become more acidic due to sulfur dioxide emissions, forming "acid rain".

(a) Construct a balanced equation, including state symbols, for the first dissociation of carbonic acid, H₂CO₃, in water.
[1]
........................................................................................................................................

(b) Write the expression for the acid dissociation constant, $K_a$ , for this reaction.
[1]
........................................................................................................................................

(c) Sulfur dioxide, SO₂, dissolves in water to form sulfurous acid, H₂SO₃.
The $K_a$ of H₂SO₃ is significantly larger than the $K_a$ of H₂CO₃.
Explain what this indicates about the relative strengths of these two acids.
[1]
........................................................................................................................................
........................................................................................................................................

(d) Acid rain can damage limestone buildings (calcium carbonate). Write an ionic equation for the reaction between calcium carbonate and hydrogen ions.
[1]
........................................................................................................................................

7. Enzyme activity is highly dependent on pH. The enzyme pepsin, found in the stomach, works optimally at pH 2.0.

(a) Calculate the concentration of hydroxide ions, [OH⁻], in a solution with pH 2.0 at 298 K. ( $K_w = 1.0 \times 10^{-14}$ mol² dm⁻⁶)
[2]
........................................................................................................................................
........................................................................................................................................
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(b) Explain why pepsin would lose its catalytic activity if the pH of the stomach were raised to pH 7.0. Refer to the structure of the enzyme in your answer.
[2]
........................................................................................................................................
........................................................................................................................................
........................................................................................................................................

8. Propanoic acid, CH₃CH₂COOH, is a weak acid with $K_a = 1.3 \times 10^{-5}$ mol dm⁻³.

(a) Calculate the pH of a 0.10 mol dm⁻³ solution of propanoic acid. State any assumptions made.
[3]
........................................................................................................................................
........................................................................................................................................
........................................................................................................................................
........................................................................................................................................
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(b) This solution is diluted by adding an equal volume of water.
(i) State the effect of this dilution on the value of $K_a$ .
[1]
........................................................................................................................................

(ii) State the effect of this dilution on the pH of the solution.
[1]
........................................................................................................................................

9. Ammonia, NH₃, is a weak base.

(a) Write an equation for the reaction of ammonia with water.
[1]
........................................................................................................................................

(b) Ammonium chloride, NH₄Cl, is a salt derived from ammonia and hydrochloric acid.
Predict whether an aqueous solution of ammonium chloride will be acidic, alkaline, or neutral. Explain your answer with reference to the ions present.
[2]
........................................................................................................................................
........................................................................................................................................
........................................................................................................................................

10. A student is asked to distinguish between two white solids: sodium chloride and ammonium chloride.

(a) Describe a simple chemical test, including the expected observations, that would distinguish between these two solids.
[2]
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........................................................................................................................................
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(b) Explain the difference in thermal stability between sodium chloride and ammonium chloride.
[2]
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Section C: Extended Response and Synthesis

11. Lactic acid (2-hydroxypropanoic acid) is produced in muscles during intense exercise and in milk during fermentation.

(a) Lactic acid contains two functional groups. Identify them.
[1]
........................................................................................................................................

(b) Lactic acid reacts with sodium carbonate. Write a balanced equation for this reaction.
[2]
........................................................................................................................................
........................................................................................................................................

(c) In a fermentation process, the buildup of lactic acid can inhibit the enzymes responsible for the reaction.
(i) Explain why the accumulation of acid inhibits enzyme activity.
[2]
........................................................................................................................................
........................................................................................................................................

(ii) Calcium hydroxide is sometimes added to fermentation tanks. Explain the purpose of adding calcium hydroxide in this context.
[1]
........................................................................................................................................

12. The table below shows the pH values of 0.1 mol dm⁻³ solutions of three different acids.

Acid	Formula	pH
Hydrochloric acid	HCl	1.0
Ethanoic acid	CH₃COOH	2.9
Chloroethanoic acid	CH₂ClCOOH	1.9

(a) Explain why hydrochloric acid has a lower pH than ethanoic acid.
[2]
........................................................................................................................................
........................................................................................................................................

(b) Explain why chloroethanoic acid is a stronger acid than ethanoic acid. Refer to the electronic effects of the chlorine atom in your answer.
[3]
........................................................................................................................................
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........................................................................................................................................
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13. Magnesium oxide and silicon(IV) oxide are both oxides of Period 3 elements.

(a) Describe the structure and bonding in magnesium oxide.
[2]
........................................................................................................................................
........................................................................................................................................

(b) Describe the structure and bonding in silicon(IV) oxide.
[2]
........................................................................................................................................
........................................................................................................................................

(c) Explain why magnesium oxide has a much higher melting point than silicon(IV) oxide, or vice versa? (Note: Check data booklet values if unsure, but explain based on bonding strength).
Correction for H1 Level: Actually, SiO₂ (mp ~1700°C) and MgO (mp ~2800°C). MgO is higher.
Explain why magnesium oxide has a higher melting point than silicon(IV) oxide.
[2]
........................................................................................................................................
........................................................................................................................................

14. A buffer solution is made by mixing 50 cm³ of 0.10 mol dm⁻³ ethanoic acid with 50 cm³ of 0.10 mol dm⁻³ sodium ethanoate.

(a) Calculate the moles of ethanoic acid and ethanoate ions in the mixture.
[1]
........................................................................................................................................

(b) 10 cm³ of 0.10 mol dm⁻³ HCl is added to this buffer.
Calculate the new pH of the solution. ( $K_a$ for ethanoic acid = $1.7 \times 10^{-5}$ mol dm⁻³)
[4]
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........................................................................................................................................
........................................................................................................................................
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15. Solubility equilibria are important in qualitative analysis.

(a) Silver chloride, AgCl, is sparingly soluble in water.
Write the equilibrium equation for the dissolution of AgCl.
[1]
........................................................................................................................................

(b) Explain why AgCl dissolves in dilute aqueous ammonia.
[2]
........................................................................................................................................
........................................................................................................................................

(c) Why does AgI not dissolve in dilute aqueous ammonia?
[1]
........................................................................................................................................

16. Consider the reaction between ammonia gas and hydrogen chloride gas.

(a) Write a balanced equation for this reaction.
[1]
........................................................................................................................................

(b) Identify the type of bonding formed in the product.
[1]
........................................................................................................................................

(c) The product is soluble in water. Predict the pH of the resulting solution and explain why.
[2]
........................................................................................................................................
........................................................................................................................................

17. Titration curves provide information about acid-base strength.

(a) Sketch the pH curve for the titration of 25 cm³ of 0.1 mol dm⁻³ ammonia (weak base) with 0.1 mol dm⁻³ hydrochloric acid (strong acid).
Label the equivalence point and the buffer region.
[3]

pH
14 |
   |
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   |
   |
   |
   |
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   |
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 0 |________________________________________ Volume of HCl / cm³
    0                                    50

(b) Suggest a suitable indicator for this titration and explain your choice.
[2]
........................................................................................................................................
........................................................................................................................................

18. Hydrofluoric acid, HF, is a weak acid with $K_a = 5.6 \times 10^{-4}$ mol dm⁻³.

(a) Calculate the pH of a 0.010 mol dm⁻³ solution of HF.
[3]
........................................................................................................................................
........................................................................................................................................
........................................................................................................................................
........................................................................................................................................

(b) Unlike other hydrogen halides, HF is a weak acid. Suggest a reason for this, referring to bond strength.
[1]
........................................................................................................................................

19. The following salts are dissolved in water:
NaCl, NH₄NO₃, CH₃COONa.

(a) Which salt will produce a neutral solution? Explain.
[1]
........................................................................................................................................

(b) Which salt will produce an alkaline solution? Explain with an equation.
[2]
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(c) Which salt will produce an acidic solution? Explain with an equation.
[2]
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20. An unknown acid, HA, has a concentration of 0.10 mol dm⁻³ and a pH of 3.0.

(a) Calculate the value of $K_a$ for this acid.
[3]
........................................................................................................................................
........................................................................................................................................
........................................................................................................................................
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(b) Is this acid stronger or weaker than ethanoic acid ( $K_a = 1.7 \times 10^{-5}$ )? Explain.
[1]
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End of Paper

Answers

TuitionGoWhere Exam Practice (AI) - Answer Key

A-Level Chemistry H1 Practice Paper - Version 3

Topic: Acids, Bases and Salts

Total Marks: 40

Section A: Structured Questions

1.
(a) A weak acid is an acid that partially dissociates (or ionizes) in water. [1]
(b) CH₃COOH(aq) ⇌ CH₃COO⁻(aq) + H⁺(aq)
Marking: Reversible arrow (⇌) required. State symbols required. [1]
(c) Ethanoic acid can form dimers via hydrogen bonding between two molecules (two H-bonds per dimer). Ethanol forms hydrogen bonds but not stable dimers to the same extent. More energy is required to break the intermolecular forces in ethanoic acid. [2]
Note: Accept reference to stronger intermolecular forces due to dimerization.

2.
(a) $[H^+] = 10^{-pH} = 10^{-2.90} = 1.26 \times 10^{-3}$ mol dm⁻³. [1]
(b) Moles of NaOH = $0.100 \times \frac{20.0}{1000} = 0.0020$ mol.
Ratio of C₆H₅COOH : NaOH is 1:1.
Moles of acid = 0.0020 mol.
Concentration of acid = $\frac{0.0020}{\frac{25.0}{1000}} = 0.080$ mol dm⁻³. [2]
(c) Sketch:

Start pH around 3 (weak acid).
Gradual rise (buffer region).
Vertical jump at equivalence point (20 cm³).
Equivalence point pH > 7 (basic salt).
Levels off at high pH (excess strong base).
Marks: Shape [1], Equivalence point labeled correctly at pH > 7 and Vol = 20 [1]. [2]

3.
(a) A solution that resists changes in pH when small amounts of acid or base are added. [1]
(b) $pH = pK_a + \log_{10} \left( \frac{[salt]}{[acid]} \right)$
Since [salt] = [acid], $\log(1) = 0$ .
$pH = pK_a = -\log_{10}(1.7 \times 10^{-5}) = 4.77$ . [2]
(c) The added H⁺ ions react with the ethanoate ions (CH₃COO⁻) from the salt:
CH₃COO⁻(aq) + H⁺(aq) → CH₃COOH(aq).
This removes most of the added H⁺, keeping the pH relatively constant. [2]

4.
(a) Amphoteric substances can act as both an acid and a base. [1]
(b) (i) Al₂O₃(s) + 6H⁺(aq) → 2Al³⁺(aq) + 3H₂O(l) [1]
(ii) Al₂O₃(s) + 2OH⁻(aq) + 3H₂O(l) → 2[Al(OH)₄]⁻(aq)
Accept: Al₂O₃ + 2NaOH + 3H₂O → 2NaAl(OH)₄ [1]

5.
(a) $K_{sp} = [Mg^{2+}][OH^-]^2$ [1]
(b) Let solubility be $s$ mol dm⁻³.
$[Mg^{2+}] = s$ , $[OH^-] = 2s$ .
$K_{sp} = (s)(2s)^2 = 4s^3$ .
$2.0 \times 10^{-11} = 4s^3$ .
$s^3 = 5.0 \times 10^{-12}$ .
$s = \sqrt[3]{5.0 \times 10^{-12}} = 1.71 \times 10^{-4}$ mol dm⁻³. [2]
(c) Common ion effect. MgCl₂ provides Mg²⁺ ions.
According to Le Chatelier’s principle, increasing [Mg²⁺] shifts the equilibrium $Mg(OH)_2(s) \rightleftharpoons Mg^{2+}(aq) + 2OH^-(aq)$ to the left, causing precipitation and decreasing solubility. [2]

Section B: Data Interpretation and Application

6.
(a) H₂CO₃(aq) ⇌ H⁺(aq) + HCO₃⁻(aq) [1]
(b) $K_a = \frac{[H^+][HCO_3^-]}{[H_2CO_3]}$ [1]
(c) H₂SO₃ has a larger $K_a$ , meaning it dissociates to a greater extent than H₂CO₃. Therefore, H₂SO₃ is a stronger acid. [1]
(d) CaCO₃(s) + 2H⁺(aq) → Ca²⁺(aq) + H₂O(l) + CO₂(g) [1]

7.
(a) $[H^+] = 10^{-2} = 0.01$ mol dm⁻³.
$[OH^-] = \frac{K_w}{[H^+]} = \frac{1.0 \times 10^{-14}}{0.01} = 1.0 \times 10^{-12}$ mol dm⁻³. [2]
(b) At pH 7, the concentration of H⁺ is much lower than optimal. The change in pH disrupts the ionic and hydrogen bonds maintaining the tertiary structure of the enzyme. This causes denaturation, changing the shape of the active site so the substrate can no longer bind. [2]

8.
(a) $K_a = \frac{[H^+]^2}{[HA]}$ .
Assumption: $[H^+]_{eq} \approx [H^+]_{initial}$ from dissociation, and $[HA]_{eq} \approx [HA]_{initial}$ (dissociation is small).
$[H^+] = \sqrt{K_a \times [HA]} = \sqrt{1.3 \times 10^{-5} \times 0.10} = \sqrt{1.3 \times 10^{-6}} = 1.14 \times 10^{-3}$ mol dm⁻³.
$pH = -\log(1.14 \times 10^{-3}) = 2.94$ . [3]
(b) (i) No change. $K_a$ is a constant at a given temperature. [1]
(ii) pH increases (becomes less acidic) because [H⁺] decreases upon dilution. [1]

9.
(a) NH₃(aq) + H₂O(l) ⇌ NH₄⁺(aq) + OH⁻(aq) [1]
(b) Acidic. NH₄Cl dissociates into NH₄⁺ and Cl⁻.
NH₄⁺ is the conjugate acid of a weak base (NH₃) and hydrolyzes:
NH₄⁺(aq) + H₂O(l) ⇌ NH₃(aq) + H₃O⁺(aq), producing H⁺ ions.
Cl⁻ is the conjugate base of a strong acid and does not hydrolyze. [2]

10.
(a) Add aqueous NaOH and warm.
Ammonium chloride produces ammonia gas (pungent smell, turns damp red litmus blue).
Sodium chloride shows no reaction/no gas. [2]
(b) Ammonium chloride is thermally unstable and sublimes/decomposes on heating:
NH₄Cl(s) ⇌ NH₃(g) + HCl(g).
Sodium chloride has a giant ionic lattice with strong electrostatic forces, requiring very high temperatures to melt; it does not decompose easily. [2]

11.
(a) Hydroxyl group (-OH) and Carboxyl group (-COOH). [1]
(b) 2CH₃CH(OH)COOH + Na₂CO₃ → 2CH₃CH(OH)COONa + H₂O + CO₂ [2]
(c) (i) Accumulation of acid lowers pH. This disrupts the tertiary structure of the enzyme (denaturation) by breaking ionic/hydrogen bonds, altering the active site shape. [2]
(ii) Calcium hydroxide is a base. It neutralizes the lactic acid, preventing the pH from dropping too low and inhibiting the enzymes. [1]

12.
(a) HCl is a strong acid and fully dissociates, giving a high [H⁺]. Ethanoic acid is weak and partially dissociates, giving a lower [H⁺] and thus higher pH. [2]
(b) The chlorine atom is electronegative. It exerts an electron-withdrawing inductive effect (-I effect). This withdraws electron density from the O-H bond in the carboxyl group, weakening it and making the H⁺ ion easier to lose. It also stabilizes the resulting carboxylate anion by dispersing the negative charge. [3]

13.
(a) Giant ionic lattice. Strong electrostatic forces of attraction between Mg²⁺ and O²⁻ ions. [2]
(b) Giant covalent (macromolecular) structure. Strong covalent bonds between Si and O atoms throughout the lattice. [2]
(c) MgO has a higher melting point. The electrostatic forces in the ionic lattice of MgO (charges +2/-2) are stronger than the covalent bonds in SiO₂?
Correction/Refinement for H1: Actually, comparing lattice energy vs bond energy is complex. Standard A-Level answer: MgO has very high lattice energy due to +2/-2 charges. SiO₂ has strong covalent bonds. MgO mp ~2800°C, SiO₂ ~1700°C.
Explanation: The electrostatic attraction between Mg²⁺ and O²⁻ is extremely strong due to the high charge density, requiring more energy to overcome than the covalent network breaking in SiO₂. [2]

14.
(a) Moles acid = $0.10 \times \frac{50}{1000} = 0.005$ mol.
Moles salt = $0.10 \times \frac{50}{1000} = 0.005$ mol. [1]
(b) Moles HCl added = $0.10 \times \frac{10}{1000} = 0.001$ mol.
H⁺ reacts with CH₃COO⁻:
New moles CH₃COO⁻ = $0.005 - 0.001 = 0.004$ mol.
New moles CH₃COOH = $0.005 + 0.001 = 0.006$ mol.
Total volume = 110 cm³ (cancels out in ratio).
$pH = pK_a + \log \left( \frac{0.004}{0.006} \right)$ .
$pK_a = 4.77$ .
$pH = 4.77 + \log(0.667) = 4.77 - 0.176 = 4.59$ . [4]

15.
(a) AgCl(s) ⇌ Ag⁺(aq) + Cl⁻(aq) [1]
(b) Ag⁺ ions react with NH₃ to form the complex ion [Ag(NH₃)₂]⁺. This reduces [Ag⁺], shifting the solubility equilibrium to the right, dissolving more AgCl. [2]
(c) AgI has a much smaller $K_{sp}$ (is much less soluble) than AgCl. The equilibrium concentration of Ag⁺ is too low to form the complex ion with dilute ammonia effectively. [1]

16.
(a) NH₃(g) + HCl(g) → NH₄Cl(s) [1]
(b) Ionic bonding (in the solid lattice) / Coordinate covalent (dative) bond formation within the ammonium ion.
Accept: Ionic. [1]
(c) Acidic. NH₄Cl is a salt of a weak base and strong acid. NH₄⁺ hydrolyzes to produce H⁺. [2]

17.
(a) Sketch:

Start pH ~11 (weak base).
Gradual drop.
Vertical drop at equivalence point.
Equivalence point pH < 7 (acidic salt).
Levels off at low pH.
Marks: Shape [1], Equivalence point pH < 7 [1], Buffer region indicated [1]. [3]
(b) Methyl orange. The equivalence point is in the acidic range (pH 3-5). Methyl orange changes color in this range (3.1-4.4). Phenolphthalein changes in basic range, so it would change color before the equivalence point. [2]

18.
(a) $[H^+] = \sqrt{K_a \times [HA]} = \sqrt{5.6 \times 10^{-4} \times 0.010} = \sqrt{5.6 \times 10^{-6}} = 2.37 \times 10^{-3}$ .
$pH = -\log(2.37 \times 10^{-3}) = 2.63$ . [3]
(b) The H-F bond is very strong (short bond length, high bond energy) due to the small size of F and good orbital overlap. This makes it difficult for the H⁺ to dissociate. [1]

19.
(a) NaCl. Derived from strong acid (HCl) and strong base (NaOH). Neither ion hydrolyzes. [1]
(b) CH₃COONa. CH₃COO⁻ hydrolyzes: CH₃COO⁻ + H₂O ⇌ CH₃COOH + OH⁻. Produces OH⁻, making solution alkaline. [2]
(c) NH₄NO₃. NH₄⁺ hydrolyzes: NH₄⁺ + H₂O ⇌ NH₃ + H₃O⁺. Produces H⁺, making solution acidic. [2]

20.
(a) $[H^+] = 10^{-3} = 0.001$ mol dm⁻³.
$K_a = \frac{[H^+]^2}{[HA]} = \frac{(0.001)^2}{0.10} = \frac{10^{-6}}{10^{-1}} = 1.0 \times 10^{-5}$ mol dm⁻³. [3]
(b) Weaker. $K_a$ of HA ( $1.0 \times 10^{-5}$ ) is smaller than $K_a$ of ethanoic acid ( $1.7 \times 10^{-5}$ ). Smaller $K_a$ means less dissociation. [1]