R3.1 Proton transfer reactions - IB Chemistry Question Bank

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R3.1 Proton transfer reactions

Practice exam-style IB Chemistry questions for Proton transfer reactions, aligned with the syllabus and grouped by topic.

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Question 1

SL • Paper 1A

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SL • Paper 1A

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In the reaction below, the Brønsted–Lowry acid and base are:

$NH_3(g) + HCl(g) \to NH_4Cl(s)$

A.

Acid: $NH_3$ ; base: $HCl$

B.

Acid: $HCl$ ; base: $NH_3$

C.

Acid: $NH_4^+$ ; base: $Cl^-$

D.

Acid: $Cl^-$ ; base: $NH_4^+$

Question 2

SL • Paper 1A

Easy

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SL • Paper 1A

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The conjugate base of $HSO_4^-$ is

A.

$SO_3^{2-}$

B.

$SO_4^{2-}$

C.

$HSO_4^{2-}$

D.

$H_2SO_4$

Question 3

HL • Paper 1A

Easy

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HL • Paper 1A

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The strongest weak acid in this set is

A.

acid with $pK_a=2.10$

B.

acid with $pK_a=9.25$

C.

acid with $pK_a=4.76$

D.

acid with $pK_a=6.35$

Question 4

SL • Paper 1A

Easy

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SL • Paper 1A

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The species that is amphiprotic is

A.

$CO_3^{2-}$

B.

$NH_4^+$

C.

$Cl^-$

D.

$HCO_3^-$

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Question 5

SL • Paper 1A

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SL • Paper 1A

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The $pH$ of a solution with $[H^+] = 4.0 \times 10^{-3}\ \text{mol dm}^{-3}$ is

A.

$11.60$

B.

$2.40$

C.

$0.40$

D.

$3.40$

Question 6

SL • Paper 1A

Easy

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SL • Paper 1A

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At $298\ \text{K}$ , a solution has $[H^+] = 2.0 \times 10^{-5}\ \text{mol dm}^{-3}$ and $[OH^-] = 5.0 \times 10^{-10}\ \text{mol dm}^{-3}$ . Its acid-base character is

A.

neutral, because $[H^+][OH^-]=K_w$

B.

acidic, because $[H^+] > [OH^-]$

C.

basic, because $[OH^-] < 10^{-7}\ \text{mol dm}^{-3}$

D.

basic, because $[H^+] > [OH^-]$

Question 7

SL • Paper 1A

Easy

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SL • Paper 1A

Easy

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Hydrochloric acid is titrated with aqueous sodium hydroxide. The correct general pH curve is

A.

B.

C.

D. Question 8

HL • Paper 1A

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HL • Paper 1A

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At $298\ \text{K}$ , the $pH$ of a solution with $[OH^-] = 2.5 \times 10^{-3}\ \text{mol dm}^{-3}$ is

A.

$11.40$

B.

$2.60$

C.

$11.00$

D.

$3.00$

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Question 9

SL • Paper 2

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SL • Paper 2

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Ammonia dissolves in water according to the following equilibrium.

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

Deduce the Brønsted–Lowry acid and base on the reactant side.

[2]

State one conjugate acid–base pair in this reaction.

[1]

Question 10

SL • Paper 2

Easy

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SL • Paper 2

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Conjugate acid–base pairs differ by one proton.

Deduce the conjugate base of $H_2PO_4^-$ and the conjugate acid of $CO_3^{2-}$ .

[2]

Explain why $H_2SO_4$ and $SO_4^{2-}$ are not a conjugate acid–base pair.

[1]

Question 11

HL • Paper 1A

Medium

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HL • Paper 1A

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At $298\ \text{K}$ , $K_a$ for the weak acid $HA$ is $6.3 \times 10^{-5}$ . The $K_b$ of its conjugate base $A^-$ is

A.

$1.6 \times 10^{-10}$

B.

$1.6 \times 10^{-5}$

C.

$6.3 \times 10^{9}$

D.

$6.3 \times 10^{-19}$

Question 12

HL • Paper 1A

Medium

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HL • Paper 1A

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The hydrolysis equation responsible for the basic $pH$ of aqueous sodium ethanoate is

A.

$CH_3COO^-(aq) + H_2O(l) \rightleftharpoons CH_3COOH(aq) + OH^-(aq)$

B.

$CH_3COOH(aq) + OH^-(aq) \rightleftharpoons CH_3COO^-(aq) + H_2O(l)$

C.

$Na^+(aq) + H_2O(l) \rightleftharpoons NaOH(aq) + H^+(aq)$

D.

$CH_3COO^-(aq) + H^+(aq) \to CH_3COOH(aq)$

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Question 13

HL • Paper 1A

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HL • Paper 1A

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A weak monoprotic acid is titrated with aqueous sodium hydroxide. The correct general pH curve is

A.

B.

C.

D. Question 14

HL • Paper 1A

Medium

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HL • Paper 1A

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A buffer contains $CH_3COOH$ and $CH_3COO^-$ . The $pK_a$ of $CH_3COOH$ is $4.76$ and $[CH_3COO^-]/[CH_3COOH]=10.0$ . Its $pH$ at $298\ \text{K}$ is

A.

$5.76$

B.

$4.76$

C.

$3.76$

D.

$10.0$

Question 15

SL • Paper 2

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SL • Paper 2

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The hydrogencarbonate ion, $HCO_3^-$ , is amphiprotic in aqueous solution.

Formulate an equation showing $HCO_3^-$ acting as a Brønsted–Lowry acid in water.

[1]

Formulate an equation showing $HCO_3^-$ acting as a Brønsted–Lowry base in water.

[1]

Question 16

SL • Paper 2

Medium

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SL • Paper 2

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A sample of rainwater has a pH of 5.60 at $298\ \text{K}$ .

Calculate $[H^+]$ in the rainwater.

[2]

State whether this rainwater is acidic, neutral or basic at $298\ \text{K}$ .

[1]

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Question 17

HL • Paper 2

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HL • Paper 2

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At $298\ \text{K}$ , the hydroxide ion concentration in an alkaline cleaning solution is $3.20 \times 10^{-4}\ \text{mol dm}^{-3}$ .

Calculate the pOH of the solution.

[1]

Calculate the pH of the solution.

[1]

Question 18

HL • Paper 2

Medium

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HL • Paper 2

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The acid dissociation data for two weak acids at $298\ \text{K}$ are:

$HX$ : $pK_a = 3.18$

$HY$ : $K_a = 1.6 \times 10^{-5}\ \text{mol dm}^{-3}$

Calculate the $pK_a$ of $HY$ .

[1]

Identify the stronger acid and justify your answer.

[2]

Question 19

SL • Paper 1B

Medium

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SL • Paper 1B

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The diagram shows observations when dry hydrogen chloride and ammonia gases are allowed to mix, and an aqueous proton transfer reaction involving ethanoic acid.

Deduce the Brønsted–Lowry acid and base in the reaction between hydrogen chloride and ammonia.

[1]

State the formula of the conjugate base of ethanoic acid in the aqueous reaction.

[1]

Explain, using Brønsted–Lowry theory, why the white solid forms when the two gases meet.

[2]

Question 20

SL • Paper 2

Medium

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SL • Paper 2

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Equal volumes of $0.100\ \text{mol dm}^{-3}$ hydrochloric acid and $0.100\ \text{mol dm}^{-3}$ ethanoic acid are tested separately under identical conditions.

Distinguish between a strong acid and a weak acid in terms of ionization.

[2]

Suggest one experimental observation, other than pH, that would distinguish these two acids and explain it.

[2]

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Question 21

SL • Paper 2

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SL • Paper 2

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A $25.00\ \text{cm}^3$ sample of a monoprotic strong acid is titrated with $0.100\ \text{mol dm}^{-3}$ sodium hydroxide. The equivalence point occurs after $18.60\ \text{cm}^3$ of sodium hydroxide has been added.

State the pH at the equivalence point at $298\ \text{K}$ .

[1]

Calculate the concentration of the acid.

[2]

Explain why smaller volumes of titrant should be added near the equivalence point when collecting the pH data.

[1]

Question 22

HL • Paper 2

Medium

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HL • Paper 2

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Aqueous solutions of ammonium nitrate, $NH_4NO_3$ , and sodium carbonate, $Na_2CO_3$ , are prepared at the same concentration.

Construct an equation for the hydrolysis of the ammonium ion and predict its effect on pH.

[2]

Construct an equation for the hydrolysis of the carbonate ion and predict its effect on pH.

[2]

Question 23

HL • Paper 2

Medium

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HL • Paper 2

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A monoprotic acid is titrated with a strong base. The pH curve starts above pH 2, has a buffer region, and has an equivalence point above pH 7. The half-equivalence point occurs at pH 4.74.

Identify the type of acid present and give one feature of the curve supporting your answer.

[2]

State the $pK_a$ of the acid.

[1]

State why an indicator with a transition range centred near pH 7 would be less suitable than one changing in the basic region.

[1]

Question 24

SL • Paper 1B

Medium

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SL • Paper 1B

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A student investigated the behaviour of aqueous hydrogencarbonate ions by adding them to acidic and alkaline solutions. The measured pH changes are summarized.

Substance added	Acidic solution	Alkaline solution
$HCO_3^-$ (aq)	pH 2.0 $\to$ 3.0	pH 12.0 $\to$ 11.0
ZnO(s)	white solid dissolves	white solid dissolves

Use the data to identify one observation that shows $HCO_3^-$ can act as a Brønsted–Lowry base.

[1]

Formulate an equation showing $HCO_3^-$ acting as a Brønsted–Lowry acid in water.

[1]

Explain why $HCO_3^-$ is amphiprotic whereas zinc oxide is amphoteric but not amphiprotic.

[2]

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Question 25

SL • Paper 1B

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SL • Paper 1B

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The pH of four household solutions was estimated using universal indicator and then measured using a calibrated pH probe.

Household solution	Universal indicator colour	Estimated pH range	pH probe reading / pH
Lemon juice	red	2–3	2.40
Vinegar	orange	3–4	3.40
Tap water	green	6–7	7.00
Washing-up liquid solution	blue	9–10	9.20

Calculate $[H^+]$ for the sample with pH $3.40$ .

[1]

Two acidic samples have pH values of $2.40$ and $3.40$ . Determine the ratio of their hydrogen ion concentrations.

[1]

Describe the shape expected for a graph of pH against $[H^+]$ .

[1]

Suggest why the pH probe is more suitable than universal indicator for this investigation.

[1]

Question 26

SL • Paper 1B

Medium

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SL • Paper 1B

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At $298\ \text{K}$ , the ion product constant of water is $K_w = 1.00 \times 10^{-14}\ \text{mol}^2\ \text{dm}^{-6}$ . Some ion concentrations for dilute aqueous solutions are shown.

Solution	$[\text{H}^+]$ / mol dm $^{-3}$	$[\text{OH}^-]$ / mol dm $^{-3}$
A	$2.50 \times 10^{-5}$	—
B	$1.00 \times 10^{-7}$	$1.00 \times 10^{-7}$
C	$1.00 \times 10^{-10}$	$1.00 \times 10^{-4}$

Calculate the missing $[OH^-]$ for the solution with $[H^+] = 2.50 \times 10^{-5}\ \text{mol dm}^{-3}$ .

[1]

Classify the solution in the table for which $[H^+] < [OH^-]$ .

[1]

Explain why increasing $[H^+]$ decreases $[OH^-]$ at constant temperature.

[2]

Question 27

HL • Paper 1B

Medium

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HL • Paper 1B

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At $298\ \text{K}$ , several alkaline cleaning solutions were analysed. Some pOH, pH and hydroxide ion concentration values are shown.

Solution	$[\text{OH}^-]$ / mol dm $^{-3}$	pOH	pH
Solution A	$3.20 \times 10^{-4}$	—	—
Solution B	—	—	11.80

Calculate the pOH of a solution with $[OH^-] = 3.20 \times 10^{-4}\ \text{mol dm}^{-3}$ .

[1]

Calculate the pH of this solution at $298\ \text{K}$ .

[1]

Another solution has pH $11.80$ . Determine $[OH^-]$ at $298\ \text{K}$ .

[2]

Question 28

HL • Paper 1B

Medium

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HL • Paper 1B

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The acid dissociation data for three weak acids at $298\ \text{K}$ are shown.

Weak acid	pKa at 298 K
methanoic acid	3.75
ethanoic acid	4.76
propanoic acid	4.87

Identify the strongest weak acid in the table.

[1]

Calculate $K_a$ for ethanoic acid if $pK_a = 4.76$ .

[1]

Explain why a lower $pK_a$ corresponds to a stronger weak acid.

[1]

For the reaction between methanoic acid and ethanoate ions, state which side of the equilibrium is favoured.

[1]

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Question 29

HL • Paper 2

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HL • Paper 2

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Ethanoic acid has $K_a = 1.80 \times 10^{-5}\ \text{mol dm}^{-3}$ at $298\ \text{K}$ . A solution of sodium ethanoate has concentration $0.100\ \text{mol dm}^{-3}$ . Use $K_w = 1.00 \times 10^{-14}\ \text{mol}^2\ \text{dm}^{-6}$ .

Calculate $K_b$ for the ethanoate ion.

[1]

Estimate the pH of the sodium ethanoate solution, stating the approximation used.

[3]

Question 30

HL • Paper 2

Medium

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HL • Paper 2

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A buffer solution is prepared by mixing ethanoic acid and sodium ethanoate so that the final concentrations are $0.200\ \text{mol dm}^{-3}$ $CH_3COOH$ and $0.300\ \text{mol dm}^{-3}$ $CH_3COO^-$ . For ethanoic acid, $pK_a = 4.76$ at $298\ \text{K}$ .

Calculate the pH of the buffer.

[2]

Explain why moderate dilution with distilled water has little effect on the pH but reduces the buffer capacity.

[2]

Question 31

SL • Paper 1B

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SL • Paper 1B

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Equal volumes of $0.100\ \text{mol dm}^{-3}$ hydrochloric acid and ethanoic acid were separately reacted with excess magnesium ribbon. The conductivity and pH of each acid were also measured before the magnesium was added.

Acid	Concentration / mol dm $^{-3}$	pH	Conductivity / mS cm $^{-1}$	H $_2$ at 20 s / cm $^3$	H $_2$ at 40 s / cm $^3$	H $_2$ at 60 s / cm $^3$	H $_2$ at 80 s / cm $^3$	Final H $_2$ / cm $^3$
Hydrochloric acid	0.100	1.0	39.0	14.0	25.0	29.0	30.0	30.0
Ethanoic acid	0.100	2.9	0.8	4.0	9.0	16.0	24.0	30.0

Identify the stronger acid using the data.

[1]

Explain two pieces of evidence from the data that support your answer to (a).

[2]

Explain why the data do not show that ethanoic acid is more dilute than hydrochloric acid.

[2]

Question 32

SL • Paper 1B

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SL • Paper 1B

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A $25.00\ \text{cm}^3$ sample of hydrochloric acid was titrated with $0.100\ \text{mol dm}^{-3}$ sodium hydroxide. The pH was recorded after each addition of sodium hydroxide.

State the pH at the equivalence point shown by the curve.

[1]

The equivalence volume is $20.00\ \text{cm}^3$ . Calculate the concentration of the hydrochloric acid.

[2]

Explain the shape of the curve before and near the equivalence point.

[2]

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Question 33

HL • Paper 1B

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HL • Paper 1B

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The table gives dissociation constants for weak acids and their conjugate bases at $298\ \text{K}$ . Use $K_w = 1.00 \times 10^{-14}\ \text{mol}^2\ \text{dm}^{-6}$ .

Weak acid	Conjugate base	$K_a$ / mol dm $^{-3}$
Benzoic acid	Benzoate ion	$6.30 \times 10^{-5}$
Ethanoic acid	Ethanoate ion	$1.74 \times 10^{-5}$

Calculate $K_b$ for benzoate ions if $K_a$ for benzoic acid is $6.30 \times 10^{-5}\ \text{mol dm}^{-3}$ .

[1]

Calculate $pK_b$ for benzoate ions.

[1]

Use the expressions for $K_a$ and $K_b$ to show why $K_a \times K_b = K_w$ for a conjugate acid-base pair.

[1]

Ethanoic acid has $K_a = 1.74 \times 10^{-5}\ \text{mol dm}^{-3}$ . Compare the base strengths of ethanoate and benzoate ions.

[2]

Question 34

HL • Paper 1B

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HL • Paper 1B

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Equal concentrations of four salt solutions were prepared at $298\ \text{K}$ . Their pH values were measured after calibration of the pH probe.

Salt solution	Measured pH
ammonium chloride	5.4
sodium chloride	7.0
sodium hydrogencarbonate	8.3
sodium ethanoate	8.9

Identify the salt solution that is acidic and state the ion responsible.

[1]

Construct the hydrolysis equation for the ethanoate ion in sodium ethanoate solution.

[1]

Explain why sodium chloride solution is approximately neutral.

[1]

The hydrogencarbonate solution is slightly basic. Explain this observation in terms of the reactions of $HCO_3^-$ with water.

[1]

Question 35

SL • Paper 2

Hard

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SL • Paper 2

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The hydrogencarbonate ion can take part in several proton transfer reactions.

Reaction 1: $HNO_3(aq) + H_2O(l) \to H_3O^+(aq) + NO_3^-(aq)$

Reaction 2: $HCO_3^-(aq) + H_2O(l) \rightleftharpoons CO_3^{2-}(aq) + H_3O^+(aq)$

Reaction 3: $HCO_3^-(aq) + H_2O(l) \rightleftharpoons H_2CO_3(aq) + OH^-(aq)$

Consider reaction 1.

I.

Deduce the Brønsted-Lowry acid and base in reaction 1, giving a reason for each.

[2]

II.

State the two conjugate acid-base pairs in reaction 1.

[2]

Explain, using reactions 2 and 3, why $HCO_3^-$ is amphiprotic.

[2]

Distinguish between amphiprotic and amphoteric behaviour.

[1]

Question 36

SL • Paper 2

Hard

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SL • Paper 2

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At $298\ \text{K}$ , an aqueous sample from a lake affected by acid drainage has a measured pH of $3.82$ . The value of $K_w$ at this temperature is $1.00 \times 10^{-14}\ \text{mol}^2\ \text{dm}^{-6}$ .

Calculations for the lake-water sample are carried out at $298\ \text{K}$ .

I.

Calculate $[H^+]$ in the lake-water sample.

[1]

II.

Calculate $[OH^-]$ in the lake-water sample.

[2]

A second aqueous sample at $298\ \text{K}$ has pH $5.82$ .

I.

Compare the hydrogen ion concentration of a solution of pH $3.82$ with that of a solution of pH $5.82$ .

[2]

II.

Deduce whether the second sample is acidic, neutral or basic. Explain your answer using the ion concentrations.

[2]

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Question 37

SL • Paper 2

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SL • Paper 2

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Neutralization reactions are used to prepare salts and to treat excess acidity. Magnesium oxide, sodium carbonate and ammonia are common bases.

Formulate equations for neutralization reactions.

I.

Write a balanced equation for the reaction between sulfuric acid and magnesium oxide.

[1]

II.

Write a balanced equation for the reaction between hydrochloric acid and sodium carbonate.

[2]

Identify the parent acid and parent base of ammonium ethanoate, $CH_3COONH_4$ .

[2]

Dilute sulfuric acid reacts with magnesium oxide and also reacts with magnesium metal. Explain why one reaction is a neutralization reaction, whereas the other is also a redox reaction.

[3]

Question 38

SL • Paper 2

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SL • Paper 2

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A student monitors the acidity of a coloured fruit drink during storage. The hydrogen ion concentration in one sample is $4.0 \times 10^{-5}\ \text{mol dm}^{-3}$ .

The student can use universal indicator solution or a pH probe.

I.

Compare the suitability of universal indicator and a pH probe for measuring the pH of the coloured drink over time.

[2]

II.

Outline one procedure needed before using the pH probe to improve the reliability of its readings.

[1]

The relationship between pH and $[H^+]$ is logarithmic.

I.

Calculate the pH of the sample.

[1]

II.

Sketch the general shape of a graph of pH against $[H^+]$ for aqueous solutions.

[2]

At a temperature above $298\ \text{K}$ , pure water may have a pH below $7$ . Explain why it is still neutral.

[1]

Question 39

HL • Paper 2

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HL • Paper 2

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An acid-base indicator can be represented as the weak acid $HInd$ . The undissociated form, $HInd$ , is yellow and the deprotonated form, $Ind^-$ , is red. The indicator has $pK_a = 5.1$ .

The colour of the indicator depends on the position of its equilibrium.

I.

Write the equilibrium equation and the expression for $K_{Ind}$ .

[2]

II.

Explain why the indicator is yellow in strongly acidic solution and red in alkaline solution.

[2]

Indicators are used to signal the end point in titrations.

I.

Distinguish between the end point and the equivalence point in an acid-base titration.

[2]

II.

Explain why universal indicator is not suitable for obtaining a precise titration end point.

[1]

Question 40

HL • Paper 1B

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HL • Paper 1B

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A weak monoprotic acid was titrated with sodium hydroxide. For comparison, other possible acid-base titration curves and several indicator ranges are shown.

Item	NaOH added / cm $^3$	Weak acid pH	Strong acid pH	Indicator / transition range
Titration	0.0	2.90	1.00
Titration	5.0	4.20	1.18
Titration	10.0	4.62	1.37
Titration	12.5	4.80	1.48
Titration	20.0	5.40	1.95
Titration	24.5	6.49	2.99
Titration	25.0	8.75	7.00
Titration	25.5	11.00	11.00
Indicator				methyl orange, 3.1-4.4
Indicator				methyl red, 4.2-6.3
Indicator				bromothymol blue, 6.0-7.6
Indicator				phenolphthalein, 8.3-10.0

Identify the feature of the curve that shows the acid being titrated is weak.

[1]

The half-equivalence point occurs at pH $4.80$ . State the $pK_a$ of the acid.

[1]

Select a suitable indicator from the table for this titration and justify your choice.

[2]

Suggest how the titration data should be collected near the equivalence point to improve the curve.

[1]

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Question 41

HL • Paper 1B

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HL • Paper 1B

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Several ethanoic acid/ethanoate buffer mixtures were prepared at $298\ \text{K}$ . The table shows their composition and the pH measured before and after dilution or addition of acid. For ethanoic acid, $pK_a = 4.76$ .

Buffer	$CH_3COOH$ / mol $dm^{-3}$	$CH_3COO^-$ / mol $dm^{-3}$	Treatment	Initial pH	Final pH
A	0.100	0.100	10-fold dilution	4.76	4.75
B	0.100	0.100	0.010 mol HCl added to 1.00 $\text{dm}^{3}$ buffer	4.76	4.67
C	0.400	0.400	0.010 mol HCl added to 1.00 $\text{dm}^{3}$ buffer	4.76	4.74

Calculate the pH of a buffer in which $[CH_3COO^-]/[CH_3COOH] = 0.500$ .

[1]

Explain why dilution by the same factor changes the buffer pH only slightly.

[1]

Write the equation for the reaction that minimizes pH change when a small amount of strong acid is added to this buffer.

[1]

Use the data to evaluate which of two buffers with the same concentration ratio has the greater buffer capacity.

[2]

Question 42

SL • Paper 2

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SL • Paper 2

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Equal volumes of $0.100\ \text{mol dm}^{-3}$ hydrochloric acid and $0.100\ \text{mol dm}^{-3}$ ethanoic acid are separately reacted with excess calcium carbonate at the same temperature. The carbon dioxide produced is collected over time.

Consider the acid strength of hydrochloric acid and ethanoic acid.

I.

Distinguish between a strong acid and a weak acid in terms of ionization.

[2]

II.

Write equations to show the ionization of hydrochloric acid and ethanoic acid in water, using appropriate arrows.

[2]

Explain how pH, electrical conductivity and the gas collection data could be used to distinguish the two acids.

[3]

Explain why hydrogen iodide is a stronger acid than hydrogen fluoride.

[1]

Question 43

SL • Paper 2

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SL • Paper 2

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A $25.00\ \text{cm}^3$ sample of hydrochloric acid is titrated with $0.1000\ \text{mol dm}^{-3}$ sodium hydroxide at $298\ \text{K}$ . The equivalence volume is $21.40\ \text{cm}^3$ .

Use the equivalence volume to determine the concentration of the hydrochloric acid.

I.

Calculate the amount, in mol, of sodium hydroxide added at the equivalence point.

[1]

II.

Calculate the concentration of the hydrochloric acid.

[2]

The titration is between a monoprotic strong acid and a strong base.

I.

Explain the meaning of the equivalence point and why its pH is $7$ in this titration at $298\ \text{K}$ .

[2]

II.

Draw the expected pH curve for adding sodium hydroxide to the hydrochloric acid. Label the equivalence point.

[3]

Question 44

HL • Paper 2

Hard

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HL • Paper 2

Hard

Calculator Permitted

Methylamine, $CH_3NH_2$ , is a weak base. At $298\ \text{K}$ , $K_b$ for methylamine is $4.4 \times 10^{-4}\ \text{mol dm}^{-3}$ . A solution has an initial methylamine concentration of $0.150\ \text{mol dm}^{-3}$ .

The ionization of methylamine in water is considered using the weak-base approximation.

I.

Write an equation for the reaction of methylamine with water and the expression for $K_b$ .

[2]

II.

Calculate $[OH^-]$ in the solution, stating the approximation used.

[2]

III.

Calculate the pH of the solution.

[2]

Evaluate the validity of the approximation used in (a)(ii).

[2]

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Question 45

HL • Paper 2

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HL • Paper 2

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At $298\ \text{K}$ , propanoic acid has $pK_a = 4.87$ . Ethanoic acid has $pK_a = 4.76$ .

Consider propanoic acid and its conjugate base, propanoate.

I.

Calculate $K_a$ for propanoic acid.

[1]

II.

Calculate $K_b$ for the propanoate ion.

[2]

The relationship between $K_a$ and $K_b$ can be derived from the expressions for a conjugate acid-base pair $HA/A^-$ .

I.

Derive the relationship $K_aK_b = K_w$ for the conjugate pair $HA/A^-$ .

[2]

II.

Compare the base strength of propanoate ions with ethanoate ions, using the $pK_a$ values given.

[2]

Question 46

HL • Paper 2

Hard

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HL • Paper 2

Hard

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The pH of a salt solution depends on whether the ions in the salt hydrolyse in water. Four salts are sodium methanoate, ammonium nitrate, sodium hydrogencarbonate and potassium chloride.

Consider the ions formed from weak parent species.

I.

Write an equation for the hydrolysis of the methanoate ion, $HCOO^-$ , and predict its effect on pH.

[2]

II.

Write an equation for the hydrolysis of the ammonium ion, $NH_4^+$ , and predict its effect on pH.

[2]

The hydrogencarbonate ion can hydrolyse in two ways.

I.

Write two equations showing $HCO_3^-$ acting as a base and as an acid in water.

[2]

II.

A solution of sodium hydrogencarbonate is slightly basic. Deduce which hydrolysis process predominates.

[1]

Predict the pH of potassium chloride solution at $298\ \text{K}$ , giving a reason.

[1]

Question 47

HL • Paper 2

Hard

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HL • Paper 2

Hard

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Two $25.00\ \text{cm}^3$ samples of monoprotic acids are titrated with sodium hydroxide of the same concentration. The pH curves are shown.

Interpret the two pH curves.

I.

Identify which curve represents a weak acid titrated with a strong base, giving two features of the curve as evidence.

[2]

II.

Explain why the equivalence point for the weak acid and strong base titration is above pH $7$ .

[2]

For curve A, the pH at the half-equivalence point is $4.88$ . The available indicators are methyl orange, pH range $3.1$ to $4.4$ , and phenolphthalein, pH range $8.2$ to $10.0$ .

I.

Determine $pK_a$ and $K_a$ for the weak acid.

[2]

II.

Choose the more suitable indicator for titration curve A and justify the choice.

[1]

Explain why smaller volumes of titrant should be added between pH measurements near the equivalence point.

[1]

Question 48

HL • Paper 2

Hard

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HL • Paper 2

Hard

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An acidic buffer is prepared at $298\ \text{K}$ by mixing $50.0\ \text{cm}^3$ of $0.200\ \text{mol dm}^{-3}$ ethanoic acid with $30.0\ \text{cm}^3$ of $0.200\ \text{mol dm}^{-3}$ sodium hydroxide. The $pK_a$ of ethanoic acid is $4.76$ .

The sodium hydroxide partially neutralizes the ethanoic acid.

I.

Calculate the amounts, in mol, of $CH_3COOH$ and $CH_3COO^-$ present after mixing.

[3]

II.

Calculate the pH of the buffer.

[2]

The buffer is tested by adding small amounts of acid and by dilution with water.

I.

Write an equation showing how the buffer resists a decrease in pH when a small amount of strong acid is added.

[1]

II.

Discuss the effect of diluting this buffer with water on its pH and its buffer capacity.

[2]

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R2.3 How far? The extent of chemical change

R3.2 Electron transfer reactions

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