IB Chemistry HL 3.1 Proton Transfer Reactions Questions

Use IB Chemistry HL questions to identify proton-transfer pairs, calculate pH and interpret neutralization or titration evidence.

Exam points

Identify Brønsted–Lowry proton donors and acceptors, conjugate acid–base pairs and amphiprotic species, then write balanced proton-transfer equations.
Calculate pH and pOH from [H₃O⁺], [OH⁻], K_w or equilibrium data, and convert between logarithmic measures with stated temperature conditions.
Use Kₐ, Kᵦ, pKₐ, pKᵦ and KₐKᵦ = K_w to compare acid/base strength, conjugate pairs and weak-acid or weak-base equilibria.
Predict salt hydrolysis and the pH of a salt solution from parent acid/base strengths, including hydrolysis-equilibrium calculations.
Analyse strong- and weak-acid/base titration curves, locate equivalence or half-equivalence regions, and select an indicator whose transition range matches the endpoint.
Write buffer reactions and determine buffer composition or pH after neutralization, using conjugate-pair ratios and the Henderson–Hasselbalch relationship.
Explain how dilution changes buffer capacity while leaving the conjugate-pair ratio, and interpret the effect of added strong acid or base.

Covered syllabus points

3.1.1—Brønsted-Lowry theory 3.1.2—Conjugate acid-base pairs 3.1.3—Amphiprotic species 3.1.4—pH scale 3.1.5—Ion product of water (Kw)3.1.6—Strong vs. weak acids/bases 3.1.7—Neutralization reactions 3.1.8—pH curves 3.1.9 (HL)—pOH scale 3.1.10 (HL)—Acid/base strength constants 3.1.11 (HL)—Conjugate pair relationship 3.1.12 (HL)—Salt hydrolysis 3.1.13 (HL)—pH curves for weak acids/bases 3.1.14 (HL)—Acid-base indicators 3.1.15 (HL)—Indicator selection 3.1.16 (HL)—Buffer solutions 3.1.17 (HL)—Buffer pH and composition

[Maximum number: 9]

Bleaches in which chlorine is the active ingredient are the most common, although some environmental groups have concerns about their use.

(a)

In aqueous chlorine the equilibrium below produces chloric(I) acid (hypochlorous acid), HOCl , the active bleach.

\mathrm{Cl}_{2}(\mathrm{aq})+\mathrm{H}_{2} \mathrm{O}(\mathrm{l}) \rightleftharpoons \mathrm{HOCl}(\mathrm{aq})+\mathrm{H}^{+}(\mathrm{aq})+\mathrm{Cl}^{-}(\mathrm{aq})

[ 9 ]

(i)

Chloric(I) acid is a weak acid, but hydrochloric acid is a strong acid. Outline how this is indicated in the equation above.

[ 1 ]

(ii)

State a balanced equation for the reaction of chloric(I) acid with water.

[ 1 ]

$\quad \mathrm{HOCl}(\mathrm{aq}) \rightleftharpoons \mathrm{H}^{+}(\mathrm{aq})+\mathrm{ClO}^{-}(\mathrm{aq}) / \mathrm{HOCl}(\mathrm{aq})+\mathrm{H}_{2}\mathrm{O}(\mathrm{l}) \rightleftharpoons \mathrm{H}_{3}\mathrm{O}^{+}(\mathrm{aq})+\mathrm{ClO}^{-}(\mathrm{aq})$ ;

(iii)

Partial neutralization of chloric(I) acid creates a buffer solution. Given that the $\mathrm{p} K_{\mathrm{a}}$ of chloric(I) acid is 7.53 , determine the pH of a solution that has $[\mathrm{HOCl}]=0.100 \mathrm{~mol} \mathrm{dm}^{-3}$ and $\left[\mathrm{ClO}^{-}\right]=0.0500 \mathrm{~mol} \mathrm{dm}{ }^{-3}$ .

[ 4 ]

$K_{\mathrm{a}}=10^{-7.53}=2.95 \times 10^{-8}\,\mathrm{mol}\,\mathrm{dm}^{-3}$ ;
$K_{\mathrm{a}}=\frac{\left[\mathrm{H}^{+}\right]\left[\mathrm{ClO}^{-}\right]}{[\mathrm{HOCl}]}=\frac{\left[\mathrm{H}^{+}\right](0.05)}{(0.1)} \approx \frac{\left[\mathrm{H}^{+}\right]}{2}=2.95 \times 10^{-8}\,\mathrm{mol}\,\mathrm{dm}^{-3}$ ;
$\left[\mathrm{H}^{+}\right]=2 \times 2.95 \times 10^{-8}=5.9 \times 10^{-8}\,\mathrm{mol}\,\mathrm{dm}^{-3}$ ;
$\mathrm{pH}=-\log\left(5.9 \times 10^{-8}\right)=7.23$ ;

(iv)

Describe, using HIn to represent the indicator in its acid form, why an indicator changes colour when excess alkali is added.

[ 3 ]

$\mathrm{HIn} \rightleftharpoons \mathrm{H}^{+}+\mathrm{In}^{-}$ ;
(weak acid in which the) acid/HIn and conjugate base/In- have different colours / OWTTE;
excess alkali shifts the equilibrium to the RHS/towards the conjugate base;

[Maximum number: 5]

Nitrous acid, $\mathrm{HNO}_{2}$ , is a weak acid which can be used to make acidic buffers.

(a)

A $1.00 \mathrm{~mol} \mathrm{dm}^{-3}$ solution of nitrous acid was prepared. This solution contains two conjugate acid-base pairs.

[ 5 ]

(i)

State the formulas of the conjugate acid and conjugate base in each pair.

Conjugate acid:
Conjugate base:
Conjugate acid:
Conjugate base:

[ 2 ]

conjugate acid $\mathrm{H}_{3} \mathrm{O}^{+}$ «（aq）» AND conjugate base $\mathrm{H}_{2} \mathrm{O}$ « $(\mathrm{l}) »$
conjugate acid $\mathrm{HNO}_{2}$ «（aq）» AND conjugate base $\mathrm{NO}_{2}{ }^{-}$ «（aq）»

(ii)

This $1.00 \mathrm{~mol} \mathrm{dm}^{-3}$ solution of nitrous acid was used to prepare a buffer with pH 3.00 .

Calculate the concentration of the conjugate base of nitrous acid required to make this buffer. The $\mathrm{pK}_{\mathrm{a}}$ of nitrous acid is 3.25.

Concentration of conjugate base:

[ 3 ]

$K_{a}<10^{-3.25}>=5.62 \times 10^{-4}$

« $\left[\mathrm{H}^{+}\right]=10^{-3.00}$ » =0.001 «mol dm ${ }^{-3}$ »
$\text { « }\left[\mathrm{A}^{-}\right]=\left(5.62 \times 10^{-4} \times 1.00\right) / 0.001 »=0.562 \text { «mol } \mathrm{dm}^{-3} \text { » }$

Alternative solution：
$\mathrm{pH}=\mathrm{p} K_{a}+\log _{10}[\text { salt }] /[\text { acid }]$ $\text { « } \log _{10}[\text { salt }] /[\text { acid }] \text { 》 }=-0.25$

OR
«[salt]/[acid]» = 0.562、

《 $\left[\mathrm{A}^{-}\right]=0.562 / 1.00 »=0.562$ «mol $\mathrm{dm}^{-3}$ » ↓

Award［3］for correct final answer．

[Maximum number: 4]

Phosphine (IUPAC name phosphane) is a hydride of phosphorus, with the formula $\mathrm{PH}_{3}$ .

(a)

Ammonia acts as a weak Brønsted-Lowry base when dissolved in water.

\mathrm{NH}_{3}(\mathrm{aq})+\mathrm{H}_{2} \mathrm{O}(\mathrm{l}) \rightleftharpoons \mathrm{NH}_{4}^{+}(\mathrm{aq})+\mathrm{OH}^{-}(\mathrm{aq})

Outline what is meant by the terms "weak" and "Brønsted-Lowry base".

Weak:

Brønsted-Lowry base:

[ 2 ]

(b)

Phosphine is usually prepared by heating white phosphorus, one of the allotropes of phosphorus, with concentrated aqueous sodium hydroxide. The equation for the reaction is:

\mathrm{P}_{4}(\mathrm{~s})+3 \mathrm{OH}^{-}(\mathrm{aq})+3 \mathrm{H}_{2} \mathrm{O}(\mathrm{l}) \rightarrow \mathrm{PH}_{3}(\mathrm{~g})+3 \mathrm{H}_{2} \mathrm{PO}_{2}^{-}(\mathrm{aq})

[ 2 ]

(i)

The ion $\mathrm{H}_{2} \mathrm{PO}_{2}^{-}$ is amphiprotic. Outline what is meant by amphiprotic, giving the formulas of both species it is converted to when it behaves in this manner.

[ 2 ]

can act as both a «Brønsted-Lowry» acid and a «Brønsted-Lowry» base OR can accept and/or donate a hydrogen ion/proton/ $\mathrm{H}^{+}$ $\mathrm{HPO}_{2}{ }^{2-}$ AND $\mathrm{H}_{3} \mathrm{PO}_{2}$

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Covered syllabus points

Question 1

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Question (iii)

Question (iv)

Question 2

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Question (ii)

Question 3

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