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39 exam-style questions with model answers, plus 52 quick multi-choice questions — every question on the site for this standard, grouped by the 13 pages of notes they come from.
Write a full answer before you reveal the model one — that comparison is where the marks come from. Every block links back to the notes that teach it.
Give the conjugate base of HF and of NH4+, and the conjugate acid of NH3 and of CH3COO−.
Explain, with equations, why a solution of ammonium chloride, NH4Cl, is acidic while a solution of sodium chloride is neutral.
Hydrogen sulfide, H2S, has Ka = 8.9 × 10−8 for its first dissociation. Sodium hydrogen sulfide, NaHS, dissolves to give the HS− ion, which is amphiprotic. Write equations showing HS− acting as both an acid and a base, identify the conjugate species in each, and explain what additional information would be needed to predict whether a solution of NaHS is acidic or basic.
Explain the difference between a strong acid and a concentrated acid, and write equations showing the behaviour of HCl and CH3COOH in water.
Equal volumes of 0.10 mol L−1 hydrochloric acid and 0.10 mol L−1 ethanoic acid are each reacted with excess magnesium. Explain how the initial rates of reaction and the total volumes of hydrogen produced compare.
A student measures the pH of 0.10 mol L−1 solutions of HCl, CH3COOH and NH4Cl as 1.00, 2.88 and 5.13 respectively. They conclude that 'HCl is the most concentrated acid and NH4Cl is the most dilute'. Evaluate this conclusion, calculate the percentage dissociation of each acid, and explain what the results reveal about the relative strengths of the three.
Calculate the pH of 0.0150 mol L−1 nitric acid, HNO3, and state whether the solution is acidic or basic.
Calculate the pH of 0.0220 mol L−1 potassium hydroxide, KOH, and explain why the calculation requires an extra step compared with finding the pH of a strong acid.
A student dilutes 10.0 mL of 0.100 mol L−1 HCl to 1.00 L and correctly finds the new pH to be 3.00. They then propose that diluting this solution by a further factor of 10 000 would give a pH of 7.00, and diluting again would make it basic. Evaluate this reasoning.
Calculate the pH of 0.200 mol L−1 hydrofluoric acid, HF, given Ka = 6.8 × 10−4.
Calculate the pH of 0.250 mol L−1 methylamine, CH3NH2, given Ka(CH3NH3+) = 2.3 × 10−11. Explain why Kb had to be derived.
A student calculates the pH of 1.0 × 10−4 mol L−1 ethanoic acid (Ka = 1.74 × 10−5) using the standard method and obtains pH 4.38. Evaluate whether the standard approximation is valid at this concentration, and explain what this reveals about the conditions under which the method may be used.
List all the species present in a 0.10 mol L−1 solution of ethanoic acid, in order of decreasing concentration. Do not include water.
List the species present in a 0.10 mol L−1 solution of ammonium chloride in order of decreasing concentration, and explain the position of each in your order.
Two solutions have the same pH of 3.0: one is 1.0 × 10−3 mol L−1 HCl and the other is a more concentrated solution of a weak acid HA. Compare and contrast the two solutions in terms of the species present, their relative concentrations, their electrical conductivity, and the volume of 0.10 mol L−1 NaOH each would require for complete neutralisation.
State what two components a buffer solution must contain, and write equations showing how a buffer made from CH3COOH and CH3COONa reacts with added H3O+ and with added OH−.
Explain why a solution containing only ethanoic acid is not a buffer, even though it contains a weak acid and some ethanoate ions from its own dissociation.
Buffer X contains 0.50 mol L−1 CH3COOH and 0.50 mol L−1 CH3COONa. Buffer Y contains 0.050 mol L−1 of each. Compare their pH values and their ability to resist pH change, and justify why a chemist choosing a buffer must consider both the ratio and the concentrations of its components.
A buffer contains 0.150 mol of methanoic acid, HCOOH, and 0.250 mol of sodium methanoate in 1.00 L of solution. Calculate its pH. Ka(HCOOH) = 1.8 × 10−4
A buffer contains 0.200 mol of CH3COOH and 0.200 mol of CH3COONa in 1.00 L. Calculate the pH before and after 0.050 mol of solid NaOH is added, assuming no volume change. Ka = 1.74 × 10−5
A student needs a buffer of pH 9.00. They have available ethanoic acid (pKa 4.76), hydrofluoric acid (pKa 3.17) and ammonium chloride (pKa of NH4+ = 9.25). Justify which system they should choose, calculate the ratio of the two components required, and explain why the other two acids would be unsuitable even though a calculation could be performed with them.
State the pH at the equivalence point for (a) a strong acid titrated with a strong base and (b) a weak acid titrated with a strong base, and explain the difference in one sentence each.
Explain why a weak acid titration curve has a buffer region, and explain how the pKa of the acid can be read directly from the curve.
A student titrates 25.0 mL of 0.100 mol L−1 HCl and separately 25.0 mL of 0.100 mol L−1 CH3COOH, each with 0.100 mol L−1 NaOH. They observe that both require exactly 25.0 mL of NaOH but that the curves differ markedly. Explain fully why the equivalence volumes are identical while the initial pH, the shape and the equivalence pH all differ. Ka(CH3COOH) = 1.74 × 10−5
State what an acid–base indicator is in chemical terms, and give the rule for selecting a suitable one for a titration.
Explain why phenolphthalein is suitable for titrating ethanoic acid with sodium hydroxide, but methyl orange is not.
A student is told that all three of methyl orange, bromothymol blue and phenolphthalein give acceptable results when titrating hydrochloric acid with sodium hydroxide, but that only one is acceptable for titrating ethanoic acid with sodium hydroxide. Justify both statements, and explain what property of the two titration curves accounts for the difference.
Write the equation for the equilibrium in a saturated solution of lead(II) iodide, PbI2, and the expression for its Ks.
Calculate the solubility of calcium fluoride, CaF2, in mol L−1, given Ks = 3.9 × 10−11, and state the concentration of each ion in the saturated solution.
A student is given Ks(AgCl) = 1.8 × 10−10 and Ks(Ag2CrO4) = 1.1 × 10−12. They conclude that silver chromate is about 160 times less soluble than silver chloride. Evaluate this conclusion, and explain the general circumstances under which Ks values may and may not be compared directly.
State what Q represents and the rule for using it to predict whether a precipitate will form.
20.0 mL of 0.0010 mol L−1 Pb(NO3)2 is mixed with 30.0 mL of 0.0020 mol L−1 KI. Predict by calculation whether a precipitate of PbI2 forms. Ks(PbI2) = 7.1 × 10−9
A student mixes 20.0 mL of 0.0188 mol L−1 AgNO3 with 30.0 mL of 0.0146 mol L−1 Al2(SO4)3 and calculates Q for Ag2SO4 as (0.0188)2(0.0146) = 5.16 × 10−6, concluding that since this is less than Ks = 1.20 × 10−5, no precipitate forms. Identify every error in this calculation, perform it correctly, and state the correct conclusion.
Explain what is meant by the common ion effect, and state its effect on the solubility of a sparingly soluble solid and on its Ks.
Calculate the solubility of Ag2CrO4 in 0.10 mol L−1 AgNO3, and compare it with its solubility in pure water. Ks(Ag2CrO4) = 1.1 × 10−12
A saturated solution of Ag2SO4 is prepared. Four solutions of equal concentration are available: HNO3, Na2SO4, NH3 and KNO3. Predict and justify the effect of adding each on the solubility of Ag2SO4, including any relevant equations.
State two ways of making a sparingly soluble solid dissolve further, and explain in one sentence why each works.
Explain, with equations, why silver chloride dissolves in aqueous ammonia but not in dilute nitric acid.
Sodium hydroxide solution is added dropwise to a solution containing Zn2+ ions. A white precipitate forms, which then dissolves as more NaOH is added. Explain both observations with equations, and explain why the same experiment with Mg2+ would give a precipitate that does not redissolve.