Equilibrium principles in aqueous systems

Practice

Recall it first, then use it — testing yourself builds the memory, where rereading the notes above only feels easier.

Recall and application shuffled together — you have to work out what each item wants before you can answer it, which is the part the exam actually tests. Both kinds feed the same spacing schedule, so anything you miss comes back sooner.

Say the answer out loud before flipping — that attempt is where the learning happens, even when you get it wrong.

Exam-style only. Instant feedback, shuffled each attempt, best score saved on this device.

What actually comes up

The shape barely changes year to year: Question One is Ks and precipitation (usually paired with a common-ion or complex-ion shift), Question Two is a buffer calculation plus identifying unlabelled solutions, Question Three is a titration curve. Grades below are read from each year's own evidence statement, and a calculation almost always earns partial credit for "correct process, minor error" even when the final number is wrong — show the substituted expression.

Across 3 years

The pattern — what to prepare

The 2025 paper ↗ Marking schedule ↗

Question One — Solubility and a complex ion

A saturated solution of iron(II) hydroxide, Fe(OH)₂, then a green precipitate of nickel(II) hydroxide, Ni(OH)₂, that dissolves when excess sodium cyanide is added.

• Fe(OH)₂ solubility = 1.01×10⁻⁵ mol L⁻¹
• Ks(Ni(OH)₂) = 6×10⁻¹⁶

• a(iii) Calculate Ks(Fe(OH)₂). Achieved
• a(iv) Dilute HCl is added to the saturated Fe(OH)₂ solution — justify, with equation(s) and no calculation, what happens to its solubility. Merit
• b(i) Explain, using equilibrium principles and a relevant equation, why the Ni(OH)₂ precipitate dissolves when excess NaCN is added. Excellence
• b(ii) 55.0 mL of 0.130 mol L⁻¹ NiSO₄ is added to 35.0 mL of KOH solution at pH 11.7. Determine whether a precipitate of Ni(OH)₂ will form. Excellence

Question Two — A methylammonium buffer

Methylammonium chloride, CH₃NH₃Cl, an acidic salt that can be mixed with methanamine, CH₃NH₂, to form a buffer.

• Ka(CH₃NH₃⁺) = 2.51×10⁻¹¹
• pKa(CH₃NH₃⁺) = 10.6

• a List all species present in a CH₃NH₃Cl solution in order of decreasing concentration (excluding water). Achieved
• b(i) Describe the function of a buffer solution, and explain the significance of CH₃NH₃⁺ + OH⁻ → CH₃NH₂ + H₂O to that function. Merit
• b(ii) 600 mL of 0.840 mol L⁻¹ CH₃NH₂; M(CH₃NH₃Cl) = 67.5 g mol⁻¹. Calculate the mass of CH₃NH₃Cl that should be added to produce a solution of pH 12.1. Excellence
• b(iii) Explain why the solution from (ii) will not function as a buffer, and outline (no calculation) how it could be modified to become one. Excellence
• c HBr, NH₄Cl and CH₃NH₂ at equal concentration are ranked in order of increasing pH — justify the order in terms of degree of dissociation and relative [H₃O⁺], with equations. Excellence

Question Three — An ammonia titration

0.0174 mol L⁻¹ HCl is added to 25.0 mL of 0.0139 mol L⁻¹ ammonia, NH₃, in a conical flask.

• Ka(NH₄⁺) = 5.75×10⁻¹⁰
• pKa(NH₄⁺) = 9.24

• a Explain, with reference to the titration curve, why methyl red (pKa 5.1) is a better equivalence-point indicator than thymol blue (pKa 1.7). Merit
• b(i) Calculate the pH at the equivalence point. Excellence
• b(ii) Compare the electrical conductivity of the flask solution before HCl is added and after 20.0 mL has been added, with relevant equation(s). Excellence
• c(ii) The 0.0174 mol L⁻¹ HCl in the burette itself has pH 1.76. Explain why the pH calculated after adding 25.0 mL of that HCl to the flask is higher than 1.76 (no calculation). Merit

The 2024 paper ↗ Marking schedule ↗

Question One — Solubility and a choice of additive

A saturated solution of silver sulfate, Ag₂SO₄, with four candidate solutions on hand: HNO₃, Na₂SO₄, NH₃, KNO₃.

• Ks(Ag₂SO₄) = 1.20×10⁻⁵

• a(iii) Calculate the solubility of Ag₂SO₄ in water at 25 °C, and give [Ag⁺] and [SO₄²⁻]. Merit
• b(i) Select and justify, with a relevant equation, which of the four listed solutions would INCREASE the solubility of Ag₂SO₄. Excellence
• b(ii) Select and justify which of the four listed solutions would DECREASE the solubility of Ag₂SO₄. Merit
• c 20.0 mL of 0.0188 mol L⁻¹ AgNO₃ is added to 30.0 mL of 0.0146 mol L⁻¹ Al₂(SO₄)₃. Predict, by calculation, whether a precipitate of Ag₂SO₄ will form. Excellence

Question Two — A sodium ethanoate buffer, and three unknowns

Sodium ethanoate, CH₃COONa, mixed with ethanoic acid, CH₃COOH, to form a buffer; separately, three unlabelled 0.110 mol L⁻¹ solutions of CH₃NH₂, CH₃COOH and NH₄Cl are told apart by conductivity and colour with bromothymol blue (pKa 7.2).

• Ka(CH₃COOH) = 1.74×10⁻⁵
• pKa(CH₃COOH) = 4.76

• a(ii) Explain how the buffer reacts when a small volume of hydrobromic acid, HBr, is added, including a balanced equation. Merit
• a(iii) 250 mL of 0.354 mol L⁻¹ CH₃COOH; M(CH₃COONa) = 82.0 g mol⁻¹. Calculate the mass of sodium ethanoate that must be added to give a buffer solution of pH 4.11. Excellence
• b Solution A: poor conductivity, yellow with bromothymol blue. B: poor, blue. C: good, yellow. Identify the three solutions, and justify the identification in terms of degree of dissociation and relative [ions], with equations (no calculation). Excellence

Question Three — An HF titration

0.169 mol L⁻¹ NaOH is added to 25.0 mL of 0.135 mol L⁻¹ hydrofluoric acid, HF.

• Ka(HF) = 6.76×10⁻⁴
• pKa(HF) = 3.17

• a(i) Show, by calculation, that the initial pH of the HF solution is 2.02. Achieved
• b(i) After a certain volume of NaOH, HF and NaF are present in a 1:9 ratio. Calculate the pH of this solution, and evaluate its buffering ability. Excellence
• b(ii) Justify, without calculation, why the pH increases rapidly between 18.0 mL and 22.0 mL of NaOH added. Merit
• c(ii) A second titration replaces HF with 0.135 mol L⁻¹ ethanoic acid, CH₃COOH (Ka = 1.74×10⁻⁵), against the same NaOH. Predict how the equivalence-point pH will compare to the HF titration, and explain your choice. Excellence

The 2023 paper ↗ Marking schedule ↗

Question One — Two solubility products

A saturated solution of lead hydroxide, Pb(OH)₂, then mixing lead nitrate with potassium iodide to test for a lead iodide, PbI₂, precipitate.

• Ks(Pb(OH)₂) = 8.00×10⁻¹⁷
• Ks(PbI₂) = 8.00×10⁻⁹

• a(iii) The saturated Pb(OH)₂ solution has pH 9.55 at 25 °C. Calculate the solubility of Pb(OH)₂. Merit
• a(iv) Explain, using equilibrium principles and a relevant equation, what happens to the solubility of Pb(OH)₂ when the pH is decreased below 3 (no calculation). Excellence
• b(i) 25.0 mL of 0.00741 mol L⁻¹ Pb(NO₃)₂ is added to 35.0 mL of 0.00613 mol L⁻¹ KI. Show, by calculation, that a precipitate of PbI₂ will form. Excellence
• b(ii) Circle sodium iodide or sodium chloride as the solution that would cause more PbI₂ to precipitate, and explain why. Excellence

Question Two — A carbonate buffer, and NH₄Cl

A buffer made from hydrogencarbonate ions, HCO₃⁻, and carbonate ions, CO₃²⁻; separately, ammonium chloride dissolved in water.

• Ka(HCO₃⁻) = 5.01×10⁻¹¹
• pKa(HCO₃⁻) = 10.3
• M(NH₄Cl) = 53.5 g mol⁻¹

• a(i) 50.0 mL of 0.211 mol L⁻¹ Na₂CO₃ mixed with 35.0 mL of 0.861 mol L⁻¹ NaHCO₃. Calculate the pH of this buffer solution. Excellence
• a(ii) A student prepares an HCO₃⁻/CO₃²⁻ solution with pH 8.50. Explain why this solution would not function effectively as a buffer. Achieved
• a(iii) Elaborate on how the ratio of HCO₃⁻:CO₃²⁻ and the pH could be adjusted to make the solution equally effective at neutralising either strong acid or strong base. Excellence
• b(ii) 500 mL of water; Ka(NH₄⁺) = 5.75×10⁻¹⁰, pKa = 9.24. Calculate the mass of NH₄Cl required to form a solution of pH 5.25. Merit
• b(iii) NH₄Cl (pH 5.61, good conductivity) and HBr (pH 1.98, good conductivity) at equal concentration. Compare the pH and electrical conductivity of the two solutions, including relevant equations. Excellence

Question Three — A methanoic acid titration

0.0893 mol L⁻¹ potassium hydroxide, KOH, is added to 15.0 mL of 0.119 mol L⁻¹ methanoic acid, HCOOH.

• Ka(HCOOH) = 1.82×10⁻⁴
• pKa(HCOOH) = 3.74

• a Calculate the pH of the methanoic acid before any KOH is added. Merit
• b(ii) Calculate the pH at the equivalence point. Excellence
• c(i) Explain, using equation(s) and no calculation, why the pH only gradually increases between about 8 mL and 12 mL of KOH added. Excellence
• c(ii) Calculate the pH in the conical flask after 23.0 mL of the KOH solution has been added. Merit

A calculation almost always earns partial credit for "correct process, minor error" even when the final number is wrong (e.g. forgetting to account for dilution) — show the substituted equilibrium expression, not just the answer.

Past exams & resources

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