Acid/Base Reactions, complete assessment covering acid-base models, strong and weak acids, pH, buffers, titrations, indicators and analysis techniques from L01-L19. 15 MC questions (auto-marked) + 5 written questions (self-marked). Complete all questions before submitting.
In the Bronsted-Lowry model, an acid is a species that:
An acid-base indicator changes colour because:
Strong acid and strong base neutralisation is represented by:
Antacids relieve excess stomach acid because they:
A weak acid is weak because it:
Two species are a conjugate acid-base pair when they differ by:
At 25 C, a solution with [H+] = 1.0 x 10-3 mol L-1 has pH:
The Ka expression for HA(aq) ↔ H+(aq) + A-(aq) is:
Weak acid neutralisation is less exothermic than strong acid neutralisation because:
A lower pKa generally means:
A buffer resists pH change best when it contains:
Concordant titres are repeated titre values that:
The best indicator for a titration has a colour-change range that:
A weak acid-strong base titration has an equivalence point:
Conductometric titration tracks equivalence by measuring changes in:
Compare Arrhenius and Bronsted-Lowry acid-base models, then identify the conjugate pairs in NH3 + H2O ↔ NH4+ + OH-.
The Arrhenius model defines acids as substances that produce H+ in water and bases as substances that produce OH- in water. The Bronsted-Lowry model is broader: acids donate protons and bases accept protons. In NH3 + H2O ↔ NH4+ + OH-, NH3 accepts a proton to become NH4+, so NH3/NH4+ is one conjugate pair. H2O donates a proton to become OH-, so H2O/OH- is the other conjugate pair.
Marks: 1, Arrhenius acid | 1, Arrhenius base | 1, Bronsted-Lowry definitions | 1, NH3/NH4+ pair | 1, H2O/OH- pairA 0.0100 mol L-1 solution of HCl is compared with a 0.0100 mol L-1 solution of ethanoic acid. Predict which has lower pH and explain using strong and weak acid behaviour.
HCl has the lower pH because it is a strong acid and is treated as fully ionised in water. For 0.0100 mol L-1 HCl, [H+] is approximately 0.0100 mol L-1, so pH = 2.00. Ethanoic acid is a weak acid, so only a fraction of its molecules ionise at equilibrium. Its [H+] is less than 0.0100 mol L-1, so its pH is higher than 2.00 even though the analytical concentration is the same.
Marks: 1, HCl lower pH | 1, strong acid full ionisation | 1, pH 2 for HCl | 1, weak acid partial ionisation | 1, higher pH explainedExplain how an ethanoic acid/ethanoate buffer responds when small amounts of acid or base are added.
An ethanoic acid/ethanoate buffer contains CH3COOH and CH3COO-. When small amounts of acid are added, CH3COO- consumes added H+ to form CH3COOH, reducing the rise in [H+]. When small amounts of base are added, CH3COOH donates H+ to neutralise OH-, forming CH3COO-. Because both conjugate partners are present in significant amounts, the equilibrium can shift either way and the pH changes only slightly until buffer capacity is exceeded.
Marks: 1, components named | 1, response to acid | 1, response to base | 1, equilibrium shift both ways | 1, buffer capacity/pH resistanceOutline how to choose an indicator for a titration and explain why phenolphthalein suits a weak acid-strong base titration better than methyl orange.
An indicator should be chosen so its transition range lies within the steep vertical region of the titration curve near the equivalence point. A weak acid-strong base titration has an equivalence point above pH 7 because the conjugate base hydrolyses water to produce OH-. Phenolphthalein changes in the alkaline range, so its colour change occurs close to this equivalence point. Methyl orange changes in the acidic range, so it would change too early and give a larger endpoint error.
Marks: 1, transition range criterion | 1, steep region/equivalence link | 1, weak acid-strong base equivalence above 7 | 1, phenolphthalein justified | 1, methyl orange limitationCompare back titration and conductometric titration, including when each is useful.
Back titration involves reacting the analyte with a known excess of reagent, then titrating the unreacted excess to determine how much reacted with the analyte. It is useful when the original reaction is slow, the sample is impure or insoluble, or a direct endpoint is difficult. Conductometric titration measures electrical conductivity as ions are consumed or formed during titration. It is useful when indicators are unsuitable, such as coloured or opaque solutions, and the equivalence point is found from the change in slope of conductivity data.
Marks: 1, back titration method | 1, back titration use | 1, conductometric method | 1, conductometric use | 1, endpoint from conductivity trend