Brønsted–Lowry Theory of Acids & Bases (Cambridge (CIE) AS Chemistry): Exam Questions

i) State what is meant by a Brønsted–Lowry base.

[1]

ii) Explain what is meant by a weak acid.

[2]

A student titrated 0.10 mol dm^-3 acid into a conical flask containing 25.0 cm³ of 0.10 mol dm^-3 of a base, recording the pH with each addition of acid. The student repeated the procedure using different combinations of acids and bases.

Fig. 5.1

Identify the curve in Fig. 5.1 that represents the titration of:

i) Ammonia and ethanoic acid.

[1]

ii) Ammonia and nitric acid.

[1]

iii) Sodium hydroxide and propanoic acid.

[1]

Identify which indicator in Table 5.2 would be most suitable for curve Z. Explain your answer.

Table 5.2

Water is amphoteric and will react with nitric acid and ammonia.

i) Construct an equation to show how water reacts with ammonia.

[1]

ii) State what is meant by the term amphoteric.

[1]

A and B are two solutions of the same concentration that have pH values of 3 and 6 respectively.

i) Identify the stronger acid. Calculate the concentration of hydrogen ions in each solution.

[2]

ii) Calculate the ratio of the hydrogen ion concentrations in A and B.

[1]

The variation of conductivity and concentration of a strong and weak monoprotic acid are shown in Fig. 2.1.

Fig. 2.1

Use the information in Fig. 2.1 to deduce which acid is strong and which is weak. Explain your reasoning.

State how the volume of 0.2 mol dm^-3 NaOH required to neutralise 20 cm³ of 0.1 mol dm^-3 acid 1 compares to the volume required to neutralise the same volume of 0.1 mol dm^-3 acid 2.

Glycolic acid, C₂H₄O₃, is an organic acid sometimes used to remove limescale, CaCO₃, from electric kettles and coffee machines.

State and explain one difference that would be observed when samples of calcium carbonate react with the same concentration of sulfuric acid and glycolic acid.

A solution of hydrochloric acid has a pH of 1 and a solution of carbonic acid has a pH of 5. Calculate the ratio of hydrogen ion concentrations of hydrochloric acid to carbonic acid.

Apart from measuring the pH, describe two methods that could be used to distinguish between equimolar solutions of carbonic acid and hydrochloric acid. State the expected observation for each method.

A student carried out two acid-base titrations. The pH ranges of four indicators are given in Table 1.1.

Table 1.1

In the first titration, 0.200 mol dm^-3 aqueous ammonia is added from a burette to 10.0 cm³ of 0.200 mol dm^-3 HCl (aq).

Sketch the pH curve for this reaction.

Using Table 1.1, explain why methyl red is a suitable indicator for the first titration but phenolphthalein is not.

In the second titration, 0.100 mol dm^-3 ethanoic acid is added from a burette to 25.0 cm³ of 0.200 mol dm^-3 NaOH (aq).

Sketch the pH curve for this reaction.

Using Table 1.1, explain why phenolphthalein is a suitable indicator for the second titration but methyl red is not.

Explain why the equivalence point of the titration in part (a) is below pH 7, and why the equivalence point of the titration in part (c) is above pH 7.

Malonic acid is a weak dicarboxylic acid with the formula C₃H₄O₄. Draw the displayed formula of malonic acid.

Ethanoic acid and malonic acid are in aqueous solution at the same concentration. State and explain which acid has the higher pH.

Apart from measuring the pH, suggest one experiment that could be used to distinguish between equimolar solutions of malonic acid and ethanoic acid. State the expected observation.

30.0 cm³ of 0.075 mol dm^-3 CH₃COOH with a pH of 2.93 is added to 10.0 cm³ of 0.150 mol dm^-3 KOH with a pH of 13.1.

On the following axes, sketch the pH curve for this addition.