R3.1 Proton transfer reactions
Practice exam-style IB Chemistry questions for Proton transfer reactions, aligned with the syllabus and grouped by topic.
Question 1
Ammonia reacts with hydrogen chloride as shown.
NH₃
A.
NH₄Cl
B.
NH₃
C.
HCl
D.
Cl⁻
Question 2
What is the conjugate base of H₂PO₄⁻?
A.
H₃PO₄
B.
HPO₄²⁻
C.
H₂PO₄²⁻
D.
PO₄³⁻
Question 3
Which species is amphiprotic?
A.
HCO₃⁻
B.
Na⁺
C.
Cl⁻
D.
CO₃²⁻
Question 4
A solution has [H⁺] = 1.0 × 10⁻³ mol dm⁻³. What is its pH?
A.
11.0
B.
3.00
C.
10.0
D.
0.001
Question 5
At 298 K, a solution has pOH = 4.20. What is its pH?
A.
18.20
B.
9.80
C.
10.20
D.
4.20
Question 6
Hydrogen sulfide reacts with water as shown.
H₂S(aq) + H₂O
⇌ HS⁻(aq) + H₃O⁺(aq)
[1]
Identify the Brønsted–Lowry acid.
[1]
Identify the Brønsted–Lowry base.
[1]
Explain why H⁺(aq) and H₃O⁺(aq) may both be used in equations for aqueous acids.
[1]
Question 7
State what is meant by a conjugate acid–base pair.
[1]
Write the formula of the conjugate acid of CO₃²⁻.
[1]
Question 8
At 298 K, a solution has [H⁺] = 2.0 × 10⁻⁵ mol dm⁻³ and [OH⁻] = 5.0 × 10⁻¹⁰ mol dm⁻³. How should the solution be classified?
A.
Basic
B.
Acidic
C.
Neutral
D.
Amphiprotic
Question 9
A 0.010 mol dm⁻³ solution of HCl and a 1.0 mol dm⁻³ solution of CH₃COOH are compared. Which statement is correct?
A.
HCl is dilute and strong; CH₃COOH is concentrated and weak.
B.
Both acids are strong because both produce H⁺ in water.
C.
HCl is dilute and weak; CH₃COOH is concentrated and strong.
D.
Both acids are weak because neither solution is pure acid.
Question 10
What are the products of the reaction between nitric acid and magnesium oxide?
A.
MgNO₃ and H₂
B.
Mg(OH)₂ and NO₂
C.
Mg(NO₃)₂ and H₂O
D.
MgCO₃ and H₂O
Question 11
Four weak acids have the following pKₐ values.
A.
Y
B.
W
C.
X
D.
Z
Question 12
At 298 K, an acid HA has Kₐ = 2.0 × 10⁻⁵. What is Kᵦ for A⁻?
A.
2.0 × 10⁻¹⁹
B.
5.0 × 10⁻¹⁰
C.
5.0 × 10⁴
D.
2.0 × 10⁹
Question 13
What is the expected pH of an aqueous solution of sodium ethanoate, CH₃COONa?
A.
Equal to 7 because all salts are neutral
B.
Below 7 because Na⁺ hydrolyses to form H₃O⁺
C.
Below 7 because CH₃COO⁻ hydrolyses to form H₃O⁺
D.
Above 7 because CH₃COO⁻ hydrolyses to form OH⁻
Question 14
For the indicator equilibrium HInd(aq) ⇌ H⁺(aq) + Ind⁻(aq), HInd is red and Ind⁻ is yellow. What colour is favoured at low pH?
A.
Yellow, because high [H⁺] shifts the equilibrium right
B.
Red, because high [H⁺] shifts the equilibrium left
C.
Red, because low [H⁺] shifts the equilibrium left
D.
Yellow, because HInd is completely ionized
Question 15
Hydrogen sulfate ions, HSO₄⁻, are amphiprotic.
Write an equation showing HSO₄⁻ acting as an acid in water.
[1]
Write an equation showing HSO₄⁻ acting as a base in water.
[1]
Distinguish between amphiprotic and amphoteric.
[1]
Question 16
A sample of rainwater has pH 4.60.
Calculate [H⁺] in mol dm⁻³.
[1]
Calculate how many times greater [H⁺] is in this rainwater than in a solution of pH 5.60.
[1]
Suggest why a pH probe is more suitable than universal indicator for comparing these two samples.
[1]
Question 17
At 298 K, Kᵥ = 1.00 × 10⁻¹⁴ mol² dm⁻⁶.
Calculate [OH⁻] in a solution where [H⁺] = 4.0 × 10⁻³ mol dm⁻³.
[1]
State whether the solution is acidic, neutral or basic.
[1]
State how [OH⁻] changes when [H⁺] is increased at constant temperature.
[1]
Question 18
At 298 K, a solution has [OH⁻] = 3.2 × 10⁻⁴ mol dm⁻³.
Calculate pOH.
[1]
Calculate pH.
[1]
Calculate [H⁺].
[1]
Question 19
Methanoic acid has Kₐ = 1.8 × 10⁻⁴ and propanoic acid has Kₐ = 1.3 × 10⁻⁵ at the same temperature.
State which acid is stronger.
[1]
Calculate pKₐ for methanoic acid.
[1]
Explain how pKₐ values compare for stronger and weaker acids.
[1]
Question 20
An indicator is represented by HInd(aq), with the equilibrium:
HInd(aq) ⇌ H⁺(aq) + Ind⁻(aq)
Write the expression for KInd.
[1]
Explain why the colour changes when alkali is added.
[1]
State why universal indicator is unsuitable for precise titration end points.
[1]
Question 21
The graph shows the pH measured for a series of aqueous solutions with different hydrogen ion concentrations.
Describe the relationship shown between pH and [H⁺].
[1]
Use the graph to estimate the pH when [H⁺] is 1.0 × 10⁻⁴ mol dm⁻³.
[1]
Calculate [H⁺] for pH 2.30.
[1]
Explain why a logarithmic scale is useful for pH.
[1]
Question 22
A student tests equal-concentration samples of acid X and acid Y. The table gives observations.
| Acid sample | pH | Conductivity / mS cm^-1 | H2 with Mg / cm^3 min^-1 |
|---|---|---|---|
| X | 2.05 | 3.9 | 11.2 |
| Y | 3.36 | 0.16 | 1.5 |
Identify which acid is stronger.
[1]
Give one piece of evidence from the table for your answer.
[1]
Explain why conductivity differs between the acids.
[1]
State why the results do not show which acid is more concentrated.
[1]
Question 23
In a titration of hydrochloric acid with sodium hydroxide at 298 K, what is the pH at the equivalence point?
A.
Greater than 7 because sodium ions are basic
B.
Equal to 7 because the salt does not hydrolyse significantly
C.
Less than 7 because acid was present initially
D.
Equal to the initial pH of the acid
Question 24
A monoprotic weak acid is titrated with a strong base. At the half-equivalence point, which relationship is correct?
A.
[HA] = 0
B.
pOH = pKₐ
C.
pH = pKₐ
D.
pH = 7.00
Question 25
A weak acid is titrated with a strong base. Which indicator range is most suitable if the equivalence point is in the basic region?
A.
pH 8.2–10.0
B.
pH 4.0–5.6
C.
pH 6.0–7.6
D.
pH 3.1–4.4
Question 26
Which mixture can act as an acidic buffer solution?
A.
CH₃COOH(aq) and CH₃COONa(aq)
B.
HNO₃(aq) and KNO₃(aq)
C.
HCl(aq) and NaCl(aq)
D.
NaOH(aq) and NaCl(aq)
Question 27
Two acids, HCl(aq) and CH₃COOH(aq), have the same concentration.
State which acid has the lower pH.
[1]
Explain the difference in pH in terms of ionization.
[1]
State one laboratory observation, other than pH, that could distinguish the two acids.
[1]
Explain why acid strength increases from HF to HI.
[1]
Question 28
Write a balanced equation for the reaction of sulfuric acid with sodium carbonate.
[1]
Write the net ionic equation for the reaction between carbonate ions and hydrogen ions.
[1]
Identify the parent acid and parent base of potassium nitrate.
[1]
State how a soluble salt can be separated from its aqueous solution after neutralization.
[1]
Question 29
A monoprotic strong acid is titrated with a strong base at 298 K.
State the pH at the equivalence point.
[1]
State what is represented by the equivalence point.
[1]
Describe two features of the pH curve as base is added.
[1]
Question 30
For the conjugate pair HA/A⁻:
HA(aq) ⇌ H⁺(aq) + A⁻(aq)
A⁻(aq) + H₂O
⇌ HA(aq) + OH⁻(aq)
[1]
Write the expression for Kₐ of HA.
[1]
Write the expression for Kᵦ of A⁻.
[1]
Show that Kₐ × Kᵦ = Kᵥ.
[1]
State the corresponding relationship between pKₐ and pKᵦ at 298 K.
[1]
Question 31
Consider aqueous solutions of NH₄Cl and Na₂CO₃.
Write an equation for hydrolysis of NH₄⁺.
[1]
Predict the effect of NH₄⁺ on pH.
[1]
Write an equation for hydrolysis of CO₃²⁻.
[1]
Predict the effect of CO₃²⁻ on pH.
[1]
Question 32
A weak base B is titrated with hydrochloric acid.
State whether the initial pH is above or below 7.
[1]
State whether the equivalence point is above, below or equal to pH 7.
[1]
State the relationship at the half-equivalence point.
[1]
Suggest why smaller volumes of acid should be added near the equivalence point.
[1]
Question 33
Distinguish between the end point and the equivalence point in a titration.
[1]
titration produces ammonium chloride at equivalence. State whether an acidic-range or basic-range indicator is more suitable.
[1]
Explain your choice in (b).
[1]
Question 34
The graph shows the titration of a monoprotic acid with sodium hydroxide at 298 K.
State the initial pH of the acid from the y-intercept.
[1]
Identify the volume of NaOH added at the equivalence point.
[1]
State whether the acid is strong or weak, using the equivalence pH.
[1]
Calculate the amount of NaOH added at equivalence if the NaOH concentration is 0.100 mol dm⁻³.
[1]
Explain why the curve becomes almost horizontal after equivalence.
[1]
Question 35
The table shows [H⁺] and [OH⁻] for three dilute aqueous solutions at 298 K.
| Solution | [H⁺] / mol dm⁻³ | [OH⁻] / mol dm⁻³ |
|---|---|---|
| A | 2.0 × 10⁻⁵ | 5.0 × 10⁻¹⁰ |
| B | 4.0 × 10⁻⁹ | x |
| C | 1.0 × 10⁻⁷ | 1.0 × 10⁻⁷ |
Classify solution A as acidic, neutral or basic.
[1]
Use Kᵥ to calculate the missing ion concentration for solution B.
[1]
Explain why [H⁺] and [OH⁻] show an inverse relationship.
[1]
Suggest how the classification of pure water may change if temperature changes.
[1]
Question 36
A student adds hydrochloric acid to sodium hydrogencarbonate and records the volume of gas produced over time.
Identify the gas produced.
[1]
Write the net ionic equation for the reaction.
[1]
Describe the change in rate of gas production shown by the graph.
[1]
Suggest why the gas production eventually stops.
[1]
State one safety or experimental precaution for this reaction.
[1]
Question 37
The table gives pH or pOH values for aqueous solutions at 298 K.
| Solution | pH | pOH | [H⁺] / mol dm⁻³ | [OH⁻] / mol dm⁻³ |
|---|---|---|---|---|
| A | 4.20 | — | — | — |
| B | — | 2.60 | — | — |
| C | 9.30 | — | — | — |
| D | 7.00 | 7.00 | 1.0 × 10⁻⁷ | 1.0 × 10⁻⁷ |
Complete the missing pOH for solution A.
[1]
Calculate [OH⁻] for solution B from its pOH.
[1]
Calculate [H⁺] for solution C from its pH.
[1]
Identify which solution is most basic.
[1]
Question 38
An acidic buffer contains 0.200 mol dm⁻³ HA and 0.300 mol dm⁻³ A⁻. The pKₐ of HA is 4.76.
Calculate the pH of the buffer.
[1]
State the effect of diluting the buffer with water on its pH.
[1]
State the effect of dilution on buffer capacity.
[1]
Question 39
The graph shows the titration of a weak monoprotic acid with sodium hydroxide.
Identify the buffer region on the curve.
[1]
Determine the equivalence volume.
[1]
Use the graph to estimate pKₐ of the acid.
[1]
Explain why the equivalence point is above pH 7.
[1]
Suggest why an indicator with a transition range around pH 5 would be unsuitable.
[1]
Question 40
A table lists indicators and their transition ranges. A titration curve for ammonia with hydrochloric acid is also shown.
State whether the equivalence point is acidic, neutral or basic.
[1]
Select the most appropriate indicator from the table.
[1]
Justify your choice using the curve and indicator range.
[1]
Distinguish between the equivalence point and the indicator end point.
[1]
Suggest why using too much indicator can introduce systematic error.
[1]
Question 41
The graph shows the pH change when small portions of HCl(aq) and NaOH(aq) are added separately to water and to a solution containing CH₃COOH/CH₃COO⁻.
Identify which solution is the buffer.
[1]
Explain how the buffer responds to added HCl.
[1]
Explain how the buffer responds to added NaOH.
[1]
State why the buffer eventually fails when enough acid is added.
[1]
Suggest why HCl/NaCl is not an effective buffer pair.
[1]
Question 42
Define Brønsted–Lowry acid and Brønsted–Lowry base.
[1]
For the reactions below, identify acids and bases and explain the proton transfers.
[1]
NH₃(aq) + H₂O
[1]
⇌ NH₄⁺(aq) + OH⁻(aq)
(ii) HNO₃(aq) + H₂O
[1]
→ H₃O⁺(aq) + NO₃⁻(aq)
[1]
Question 43
Distinguish between a strong acid and a concentrated acid.
[1]
Compare equal-concentration solutions of hydrochloric acid and ethanoic acid in terms of ionization, pH, conductivity and reaction with magnesium.
[1]
Question 44
A buffer is prepared by mixing solutions containing a weak acid HA and its salt NaA. The table gives the initial amounts and the pKₐ of HA.
| Solution mixed | Volume / cm³ | Concentration / mol dm⁻³ | pKa of HA |
|---|---|---|---|
| HA(aq) | 25.0 | 0.200 | 4.76 |
| NaA(aq) | 40.0 | 0.150 | — |
Calculate the ratio n(A⁻)/n(HA) after mixing.
[1]
Calculate the pH of the buffer.
[1]
Predict the effect on pH if the mixture is diluted tenfold with water.
[1]
Explain the effect of this dilution on buffer capacity.
[1]
Question 45
Write balanced equations for the reactions of hydrochloric acid with sodium hydroxide and with sodium carbonate.
[1]
Discuss neutralization reactions by referring to salts, net ionic equations, heat change and one method for obtaining a pure soluble salt.
[1]
Question 46
A 25.0 cm³ sample of a monoprotic strong acid is titrated with 0.100 mol dm⁻³ sodium hydroxide. The equivalence point occurs at 20.0 cm³.
Calculate the concentration of the acid.
[1]
Explain the main features of the pH curve for this titration at 298 K.
[1]
Question 47
State two features that distinguish a weak acid–strong base titration curve from a strong acid–strong base titration curve.
[1]
Evaluate how information from a weak acid–strong base pH curve can be used to determine acid strength and choose a suitable indicator.
[1]
Question 48
Describe the composition of an acidic buffer and a basic buffer.
[1]
Explain, using equations, how the ethanoic acid/ethanoate buffer resists pH change when small amounts of acid and alkali are added.
[1]
Question 49
A 0.120 mol dm⁻³ solution of a weak monoprotic acid HA has Kₐ = 6.3 × 10⁻⁵ at 298 K.
State the approximation normally used for weak acid calculations when Kₐ is small and explain why it is reasonable.
[1]
Calculate the pH of the acid solution and the Kᵦ of its conjugate base A⁻ at 298 K.
[1]
Question 50
Define hydrolysis of a salt ion and state why ions from strong acids or strong bases often have little effect on pH.
[1]
Discuss how the pH of solutions of NH₄NO₃, CH₃COONa and NaHCO₃ can be predicted from parent acid/base strengths and hydrolysis equations.
[1]
The Fast Track To Your
Best IB Coursework & College Essays
All content on this website has been developed independently from and is not endorsed by the International Baccalaureate Organization. International Baccalaureate and IB are registered trademarks owned by the International Baccalaureate Organization.