RI 15. Acid-Base Equilibria Tutorial (Questions)

-1- Raffles Institution Year 6 H2 Chemistry 2026 Tutorial 15 – Acid-Base Equilibria Self-Check Questions 1 (a) (i) Outline the Brønsted–Lowry theory of acids and bases. (ii) Identify the conjugate acid–base pairs in the following reactions: I. CH3COOH + H2SO4 ⇌ CH3COOH2+ + HSO4– II. CH3NH2(aq) + H2O(l) ⇌ CH3NH3+(aq) + OH–(aq) III. HNO2(aq) + CN–(aq) ⇌ HCN(aq) + NO2–(aq) (b) (i) Outline the Lewis theory of acids and bases. (ii) Identify the Lewis acid and base in each of the following reactions: I. (CH3)3N(g) + BF3(g) ⇌ (CH3)3NBF3(s) II. Al(OH)3(s) + OH–(aq) ⇌ [Al(OH)4]–(aq) (c) Identify the acid and base present in the following reactions, in terms of the Brønsted– Lowry theory, the Lewis theory, or both, as appropriate. (i) HBr(g) + H2O(l) → H3O+(aq) + Br–(aq) (ii) HIO + NH2– ⇌ NH3 + IO– (iii) 2 Calculate the pH of the following solutions at 298 K. (a) 0.010 mol dm−3 H2SO4 (b) 0.40 g dm−3 NaOH (c) A solution formed by mixing 14.00 cm3 of 0.10 mol dm−3 H2SO4 with 20.00 cm3 of 0.10 mol dm−3 NaOH. (d) 0.30 mol dm−3 CH3CH2COOH (pKa = 4.89) (e) 2.00 mol dm–3 CH3CH2NH2 (Kb of CH3CH2NH2 = 5.1 x 10–4 mol dm–3) [Note: CH3CH2NH2(aq) + H2O(l) ⇌ CH3CH2NH3+(aq) + OH−(aq)] (f) 0.015 mol dm−3 potassium benzoate, C6H5COO−K+ (Ka of benzoic acid, C6H5COOH = 6.5 x 10−5 mol dm−3) (g) 0.25 mol dm−3 methylammonium nitrate, CH3NH3+ NO3− (Kb of methylamine CH3NH2 = 4.4 x 10−4 mol dm−3) [Note: CH3NH3+(aq) + H2O(l) ⇌ CH3NH2(aq) + H3O+(aq)] (h) A solution containing 0.50 mol dm−3 CH3CH2COOH(aq) and 0.35 mol dm−3 CH3CH2COO−K+(aq). (Ka of CH3CH2COOH = 1.3 x 10−5 mol dm−3) (i) A solution containing 0.060 mol dm−3 of NH3 and 0.080 mol dm−3 of NH4Cl. (Kb of NH3 = 1.74 x 10−5 mol dm−3)

-2- 3 (a) A solution of 1.00 x 10–3 mol dm–3 HX has a pH of 3.25. Calculate the acid dissociation constant of HX. (b) A solution of 1.00 x 10–2 mol dm–3 sodium ethanoate has a pH of 8.37. Calculate the pKa of ethanoic acid. (c) A solution of 0.050 mol dm –3 ammonium sulfate solution has a pH of 5.12 at 298 K. Calculate the pKb of ammonia. 4 [N2009/1/6] 5 [N2008/1/11] 6 [N2010/I/12] Sulfur dioxide gas is converted to sulfate ions when it is bubbled into aqueous manganate(VII) ions. Which graph shows how the pH changes as sulfur dioxide is bubbled at a constant rate into a well -stirred solution of managate(VII) ions until its colour just fades? 7 [N2002/1/8] 8 [N2012/1/9]

-3- Practice Questions 1 A sample of apple juice has a pH of 3.50. When 25.0 cm3 of this apple juice sample is titrated against 0.100 mol dm –3 sodium hydroxide using thymolphthalein as the indicator, the average titre value is 27.50 cm3. (a) Define pH and calculate the molar concentration of hydrogen ions in apple juice. (b) Assuming that apple juice contains a single acid HA which is monobasic, calculate the molar concentration of HA in apple juice and hence explain whether HA is a strong or weak acid. (c) Calculate the degree of dissociation and the acid dissociation constant of the acid HA. (d) A sample of hydrochloric acid has the same pH as the sample of apple juice. Explain why if 1 dm3 each of apple juice and this hydrochloric acid were treated separately with the same excess amount of zinc powder, the apple juice would yield a larger volume of hydrogen gas at room temperature. 2 The pKa for ethanoic acid at 25 oC is 4.74. (a) (i) Calculate the pH of an aqueous solution of 0.50 mol dm−3 ethanoic acid. (ii) Calculate the pH of the resultant solution when equal volumes of 0.50 mol dm −3 ethanoic acid and 0.20 mol dm−3 sodium hydroxide are mixed. (b) A solution that contains only sodium ethanoate of concentration 0.050 mol dm −3 has a pH that is above 7. (i) Write an equation involving the ethanoate ion and explain why the pH of the solution is above 7. (ii) Write an expression for the base dissociation constant Kb for the ethanoate ion and calculate the value of Kb for the ethanoate ion. (iii) Hence, calculate the pH of a 0.050 mol dm−3 sodium ethanoate solution. (c) A laboratory assistant wanted to prepare a buffer solution containing ethanoic acid and sodium ethanoate. Find the volume of 0.10 mol dm–3 of sodium ethanoate that is needed to be added to 25.0 cm3 of 0.12 mol dm–3 CH3COOH(aq) to form a buffer solution of pH 5.00. 3 (a) A buffer solution is prepared by mixing 750 cm 3 of 0.20 mol dm –3 ammonium chloride solution and 500 cm3 of 0.10 mol dm–3 aqueous ammonia. Calculate the pH of this buffer solution. (Ka for the ammonium ion = 6.00 x 10–10 mol dm–3 at 25 oC) (b) Calculate the change in pH of the buffer solution prepared in (a) after the addition of (i) 0.002 mol of solid sodium hydroxide, (ii) 1.0 cm3 of 2.00 mol dm–3 hydrochloric acid. 4 Atmospheric carbon dioxide can dissolve in sea water. The three inorganic carbon-containing species present in sea water are: • Dissolved CO2(g) • CO32–(aq) • HCO3–(aq). Commented [SL1]: Adapted from 2024/P2/Q6 for new syllabus Ocean Acidification

-4- Increasing levels of CO2(g) in the atmosphere can change the proportions of these inorganic carbon-containing species present in sea water. This first happens when atmospheric CO2 dissolves in sea water: CO2(g) ⇌ CO2(aq) This can lead to ocean acidification. Reactions 4.1 and 4.2 show the two equilibria that regulate the concentration of the three inorganic carbon-containing species present in sea water and the values of their equilibrium constants. In reaction 4.1, the value for K1 assumes that the concentration of water is constant. reaction 4.1 CO2 + H2O ⇌ HCO3– + H+ K1 = 1.45 x 10−6 mol dm–3 at 298 K reaction 4.2 HCO3– ⇌ CO32– + H+ K2 = 1.10 x 10−9 mol dm–3 at 298 K When changes to the amount of dissolved carbon dioxide in the ocean occur, the pH of sea water remains approximately constant between pH 7.5 and 8.5. This is due to CO32–(aq) and HCO3–(aq) acting as a buffer. (a) Approximately 200 years ago, the average pH of sea water taken from the ocean surface was 8.2. At present, this is 8.1. Calculate the percentage increase in hydrogen ion concentration of the sea water taken from the ocean surface at present compared to 200 years ago. (b) Using information from reactions 4.1 and 4.2, suggest how the CO32–(aq) and HCO3–(aq) buffer maintains the pH in sea water as the amount of carbon dioxide gas in the atmosphere increases. (c) Deduce the most abundant inorganic carbon-containing species present in sea water at pH 8.1. Explain your answer. 5 The Ka of lactic acid is 1.4  10−4 mol dm−3 at 298 K. A buffer solution contains 1.5 mol dm−3 lactic acid and 0.5 mol dm−3 sodium lactate. Which statement correctly describes the buffer solution? A The pH of the buffer solution is 4.33. B The pH of the buffer solution decreases upon dilution with water. C The buffering capacity of the buffer solution increases upon dilution with water. D The buffer solution has a greater buffering capacity against a base compared to an acid.

-5- 6 The major acidic componen t of soured milk is lactic acid , CH3CH(OH)CO2H. When 10.0 cm3 of solution of lactic acid was titrated against 0.050 mol dm–3 sodium hydroxide the following pH readings were obtained: Volume of NaOH added /cm3 0 2 4 6 8 10 12 14 16 18 20 pH 2.5 3.1 3.4 3.7 3.9 4.1 4.4 4.7 9.1 11.6 11.8 (a) Plot a graph of these results with pH on the y –axis and volume of NaOH added on the x–axis. Identify the two major organic species present in this solution at pH 3.0 − 5.0. Explain why the slope of the graph only changes gradually in this region. (b) (i) Suggest a suitable indicator for this titration, giving a reason for your choice. (ii) For the indi