OCR A-Level Chemistry: Acids, Bases & Buffers

A buffer solution is prepared by dissolving sodium methanoate, HCOONa, in a solution of methanoic acid, HCOOH. The buffer maintains an almost constant pH when small amounts of acid or alkali are added.

Describe and explain how this buffer resists a change in pH when (i) a small amount of a strong acid is added and (ii) a small amount of a strong alkali is added. In your answer you should identify the species present in the buffer, write the relevant equilibrium, and use Le Chatelier's principle to account for what happens in each case.

(6 marks)

Propanoic acid, CH₃CH₂COOH, is a weak monobasic acid. A chemist prepares a solution of propanoic acid of concentration 0.150 mol dm⁻³ at 298 K.

Data

• Acid dissociation constant of propanoic acid: $K_a = 1.30 \times 10^{-5}$Ka​=1.30×10−5 mol dm⁻³

(a) Write the expression for the acid dissociation constant, $K_a$Ka​, of propanoic acid. (1 mark)

(b) State the two assumptions made when calculating the pH of a weak acid from its $K_a$Ka​ and concentration. (2 marks)

(c) Calculate the pH of the 0.150 mol dm⁻³ propanoic acid solution. Give your answer to two decimal places. (3 marks)

A student carried out a titration by adding 0.100 mol dm⁻³ sodium hydroxide from a burette to 25.0 cm³ of a 0.100 mol dm⁻³ solution of an acid X. The pH was recorded as the volume of sodium hydroxide added increased.

Volume of NaOH added / cm³	0.0	5.0	12.5	20.0	24.0	25.0	26.0	30.0	40.0
pH	2.9	4.1	4.8	5.4	6.1	8.7	11.1	11.9	12.3

Two indicators are available:

Indicator	pH range over which colour changes
Methyl orange	3.1 – 4.4
Phenolphthalein	8.3 – 10.0

(a) State, with a reason from the data, whether acid X is a strong acid or a weak acid. (2 marks)

(b) Determine the volume of sodium hydroxide added at the equivalence point, and state the approximate pH at that point. (1 mark)

(c) Choose the more suitable indicator for this titration and justify your choice using the data. (2 marks)

Human blood plasma is kept within a narrow pH range by a buffer system based on carbonic acid, H₂CO₃, and the hydrogencarbonate ion, HCO₃⁻. Carbonic acid acts as the weak acid and the hydrogencarbonate ion is its conjugate base:

$\text{H}_2\text{CO}_3(\text{aq}) \rightleftharpoons \text{H}^{+}(\text{aq}) + \text{HCO}_3^{-}(\text{aq})$H2​CO3​(aq)⇌H+(aq)+HCO3−​(aq)

Data

• $K_a$Ka​ of carbonic acid in plasma: $K_a = 4.30 \times 10^{-7}$Ka​=4.30×10−7 mol dm⁻³
• In healthy plasma the ratio $\dfrac{[\text{HCO}_3^{-}]}{[\text{H}_2\text{CO}_3]} = 20$[H2​CO3​][HCO3−​]​=20

(a) Calculate the pH of healthy blood plasma using these data. Give your answer to two decimal places. (3 marks)

(b) During strenuous exercise, lactic acid releases extra H⁺ into the plasma. Using the equilibrium above, explain how the hydrogencarbonate buffer responds to keep the pH almost constant, and state the effect on the ratio $\dfrac{[\text{HCO}_3^{-}]}{[\text{H}_2\text{CO}_3]}$[H2​CO3​][HCO3−​]​. (2 marks)

A technician partially neutralises ethanoic acid to make a buffer for a practical. She mixes 50.0 cm³ of 0.200 mol dm⁻³ ethanoic acid, CH₃COOH, with 20.0 cm³ of 0.200 mol dm⁻³ sodium hydroxide. The sodium hydroxide is the limiting reagent, so a buffer of ethanoic acid and sodium ethanoate is formed.

Data

• $K_a$Ka​ of ethanoic acid: $K_a = 1.74 \times 10^{-5}$Ka​=1.74×10−5 mol dm⁻³

Calculate the pH of the buffer formed. Give your answer to two decimal places.

(4 marks)

Ammonia behaves as a base in water according to the equilibrium below.

$\text{NH}_3(\text{aq}) + \text{H}_2\text{O}(\text{l}) \rightleftharpoons \text{NH}_4^{+}(\text{aq}) + \text{OH}^{-}(\text{aq})$NH3​(aq)+H2​O(l)⇌NH4+​(aq)+OH−(aq)

(a) State what is meant by a Brønsted-Lowry acid. (1 mark)

(b) Identify a conjugate acid-base pair in the equation above. (1 mark)

(c) Write the expression for the ionic product of water, $K_w$Kw​. (1 mark)