Acids, Bases & Buffers (OCR A Level Chemistry A): Exam Questions

This question is about Brønsted-Lowry acids and bases.

i) Give the meaning of the term Brønsted-Lowry base.

ii) Explain the term weak acid.

When an acid and a base react they produce a conjugate base and a conjugate acid.

acid + base  ⇋  conjugate base + conjugate acid

Write an equation to show how hydrochloric acid behaves as a strong acid when it reacts with water, and state the role of water in this reaction.

The pH of an acid or base can be calculated using the concentration of hydrogen ions. 

i) Calculate the pH of a 0.2 mol dm^-3 solution of HCl at 298K. 

ii) Calculate the hydrogen ion concentration of a solution of nitric acid which has a pH of 3. 

Acids can be classed as monobasic, dibasic and tribasic.

Give an example of a diprotic acid and explain the term dibasic.

A student measured 50.0 cm³ of 0.025 mol dm⁻³ hydrochloric acid into a beaker. The student then added 100 cm³ of distilled water into the beaker.

Calculate the pH of the solution formed after 100 cm³ of distilled water is added to the beaker. Give your answer to 3 significant figures.

At 298K, water molecules dissociate into equal quantities of ions, and the pH is 7. 

i) Write an equation to show the dissociation of water. 

ii) At 313 K, the pH of water is 6.77. Explain why water is neutral at a pH of 6.77 at this temperature.

The ionic product of water, K_w, can be used to find the pH of a strong base. Changing the temperature will affect the value for K_w.

i) Give the expression and units for the ionic product of water, K_w.

ii) As temperature is increased, the value for K_w  also increases. Explain why.

At 40 ℃ the K_w of pure water is 2.916 x 10^-14 mol² dm^-3. 

Calculate the pH of pure water. Give your answer to 3 significant figures.

Strong bases fully ionise in water, as shown by the equation of dissociation of sodium hydroxide:

NaOH (aq) →  Na⁺ (aq) + OH^- (aq)

At 298K, K_w is 1 x 10^-14 mol² dm^-6. 

Calculate the pH of a 0.05 mol dm^-3 solution of NaOH at 298 K. Give your answer to 3 significant figures.        

The ionic product of water, K_w is an equilibrium constant.

However, the expression for K_w does not include the concentration of water. Explain why.

Weak acids don’t fully ionise in solution. The equilibrium constant, K_a is used to determine the hydrogen ion concentration.

Write an expression for the acid dissociation constant, K_a for the acid HA. 

A sample of 0.01 mol dm^-3 butanoic acid has a K_a value of 1.51 x 10^-5 mol dm^-3.

i) Write an expression for the acid dissociation constant, K_a, for butanoic acid. 

ii) Calculate the pH of the 0.01 mol dm^-3 butanoic acid. Give your answer to two decimal places.

The pH of a 0.15 mol dm^-3 solution of HCN is 5.08 at 298K. 

Determine the value of K_a for HCN at 298 K. 

Give your answer to two decimal place and state its units. 

The value of K_a can be used to determine the strength of an acid. The lower the value for K_a the weaker the acid is. Hydrofluoric acid has a K_a value of 5.5 x 10^-4. In comparison, the K_a value for ethanoic acid is 1.7 x 10^-5.

Considering the expression for K_a, explain why the value of K_a is larger in hydrofluoric acid than in ethanoic acid.

Formic acid is a weak monoprotic acid. At 298K the pH of 0.025 mol dm^-3 formic acid is 5.4.

Calculate the pK_a of formic acid at 298 K. Give your answer to three significant figures.

A student titrated 0.10 mol dm^-3 acid into a conical flask containing 25.0 cm³ of 0.1 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. 

Identify which curve shown above, is the acid-base combination for the following: 

i) Ammonia and ethanoic acid.

ii) Ammonia and nitric acid.

iii) Sodium hydroxide and propanoic acid.

Identify which indicator given in the table would be most suitable for curve Y. Justify your answer. 

The pH curve shown below was obtained when a 0.150 mol dm^–3 solution of sodium hydroxide was added to 25.0 cm³ of an aqueous solution of ethanoic acid.

The half equivalence point is where half of the volume of sodium hydroxide required for neutralisation has been added to the ethanoic acid.

i) Label the graph with an X to show the position of the half equivalence point.

ii) When half of the ethanoic acid solution has been neutralised, the remaining ethanoic acid concentration is equal to that of the sodium ethanoate that had formed. The K_a =1.75 x 10^-5 mol dm^-3. Calculate the pH at this point. Give your answer to two decimal places.

The student plotted a graph of the results of the reaction between butanoic acid and potassium hydroxide. The results are shown in the pH curve below.

Use the curve to determine the value of K_a for butanoic acid at 298K. State the units for K_a. 

This question is about buffers.

Give the meaning of the term buffer solution.

Buffers can be acidic but they can also be basic. This is determined by the reactants which are combined to create the buffer

i) Describe how an acidic buffer is made.

ii) State the key feature which makes a buffer acidic. 

Ethanoic acid is a weak acid. Hydrogen carbonate ions act as a weak acid if they are in an aqueous solution.

i) Write equations for each of these weak acids at equilibrium. 

ii) A solution was made up containing sodium hydrogen carbonate and sodium carbonate. Explain how this solution would act as a buffer if a small amount of acid was added to it.

Explain how the solution of ethanoic acid works as a buffer when small amounts of alkali are added.

A buffer solution contains 0.10 mol dm^-3 concentration of methanoic acid and a 0.20 mol dm^-3 concentration of sodium methanoate. The values of K_a for methanoic acid is 1.80 × 10^–5 mol dm^–3. Calculate the pH of this buffer solution at 298 K. Give your answer to 2 decimal places.

Propanoic acid has an acid dissociation constant value, K_a, of 1.34 x 10^-5 mol dm⁻³ at 298K.

A buffer solution containing propanoic acid and sodium propanoate was prepared. 

The concentration of the propanoate ions in the buffer solution was 0.177 mol dm⁻³.  The pH of the buffer solution was 4.01. 

i) Give the expression for the acid dissociation constant K_a for propanoic acid.

ii) Calculate the concentration of the propanoic acid in the buffer solution.

Give your answer to two decimal places.

This question is about buffers.

i) Explain the term basic buffer solution.

ii) Suggest why buffers are used in the cosmetics industry. 

When we exercise, chemical changes occur which can cause our blood pH to drop. If the pH drops below 7.4 a condition called acidosis occurs. This can be very dangerous as many chemical reactions involve proteins which work at pH 7.4. The most important buffer for maintaining acid-base balance in the blood is the carbonic acid, H₂CO₃, hydrogencarbonate, HCO₃^- buffer.

i) Write a balanced equation to show the dissociation of carbonic acid, acting as a buffer solution.

ii) Explain how the carbonic acid-hydrogencarbonate buffer maintains blood pH at 7.4 when we exercise. Use your answer to part (i).

Much of the atmospheric carbon dioxide we produce is absorbed by dissolving into the oceans, forming carbonic acid. An equilibrium is set up within the oceans using the same buffer solution as in the body, described in part (b). 

In this buffer solution there is an almost unlimited amount of hydrogen carbonate ions and the carbonic acid available has a high concentration.

Using your answer to (b) part (i), suggest the changes which would lead to more CO₂ dissolving into the oceans. Explain your answer.

Lithium hydroxide monohydrate is used in the production of cathodes for lithium-ion batteries

A 4.2 g sample of lithium hydroxide monohydrate, LiOH.H₂O(s), is added to 250 cm³ of water. State the pH of the solution at 298 K. Give your answer to 1 decimal place. 

(K_w = 1.00 x 10^-14 mol² dm^-6 at 298 K)

40.00 cm³ of 0.01 mol dm^-3 hydrochloric acid, HCl (aq), is added to 60.00 cm³ of the of lithium hydroxide monohydrate solution given in part (a). State the new pH of this solution
Give your answer to 2 decimal places. 

The value of K_w is 5.476 x 10^-14 mol dm^-3 at 323K whilst at 373K the value of K_w changes to 51.3 x 10^-14.

i) Calculate the ratio of the concentrations of H⁺ ions at these two temperatures.

ii) Using your answer to part (c) (i), calculate the pH difference between water at these two temperatures.

6.00 cm³ of 0.30 mol dm^-3 potassium hydroxide solution is added to 994 cm³ of pure water at 50°C. Using the values for K_w in part (c), calculate the pH of the solution to 2 decimal places. 

This question is about acid-base reactions.

Calculate the volume of water in cm³ required to raise the pH of 35.0 cm³ of hydrochloric acid from 0.70 to 0.90 at 298 K. 

After a neutralisation reaction between sodium hydroxide and hydrochloric acid, a solution of sodium chloride is evaporated and 2.85 g of pure sodium chloride is produced. 

Determine the volume, in cm³, of 0.20 mol dm^-3 of hydrochloric acid required in the reaction to produce this mass of sodium chloride. 

44.00 cm³ of 0.20 mol dm^-3 hydrochloric acid is added to 40.00 cm³ of 0.50 mol dm^-3 sodium hydroxide. Calculate the pH of the resulting solution at 298 K. Give your answer to 2 decimal places.  

(K_w = 1.00 x 10^-14 mol² dm^-6 at 298 K)

25.00 cm³ of 0.20 mol dm^-3 barium hydroxide solution, Ba(OH)₂ (aq), is added to the 25.00 cm³ of 0.20 mol dm^-3 of hydrochloric acid. What is the pH of the new solution at 298 K? 

Give your answer to 2 decimal places.  

(K_w = 1.00 x 10^-14 mol² dm^-6 at 298 K)

Chemicals in the blood resist changes in pH when hydroxide ions enter the bloodstream.

Explain, using equations, how this occurs.

A buffer solution contains a mixture of ethanoic acid and its salt. A small amount of nitric acid is added to the buffer. 

Write an equation, including state symbols, showing how this buffer can resist the change in pH.

Determine the pH of the buffer solution which is 0.050 mol dm^-3 of ethanoic acid and 0.050 dm^-3 of sodium ethanoate. You must show working and give your answer to 2 decimal places. 

(The pK_a of ethanoic acid is 4.77)

Calculate the concentration of methanoic acid used in a buffer which has a pH of 3.45 and a methanoate salt concentration of 0.149 mol dm^-3. Give your answer to 3 significant figures.

(K_a of methanoic acid = 1.8 x 10^-4 mol dm^-3)

This question is about organic acids.

0.3825 g of an organic monobasic aromatic carboxylic acid containing only carbon, hydrogen and oxygen were dissolved in ethanol. A few drops of the indicator were added to the mixture titrated with 0.100 mol dm^-3 sodium hydroxide.

It took 25.50 cm³ of alkali to obtain the first permanent colour change.

Suggest a structure for the acid.

Table 1 shows a few common indicators and the pH ranges of their colour changes.

Suggest a suitable indicator from the table below for the procedure outlined in part (a) and explain your answer.

Table 1

The value of the acid dissociation constant, K_a, of benzoic acid is 6.3 x 10^-5 mol dm^-3 at 298K.

Calculate the pH of the 0.50 g dissolved in 250 cm³ of water. Give your answer to 2 decimal places. 

Calculate the pH of the solution formed when 0.060 g of sodium hydroxide is added to 500.00 cm³ of 0.005 mol dm^-3 of benzoic acid. Give your answer to 2 decimal places.

This question is about buffer solutions

What is the pH of a buffer solution made by dissolving 0.275 g of benzoic acid and 0.525 g of sodium benzoate in 250 cm³ of water? Give your answer to 2 decimal places. 

(K_a of benzoic acid = 6.3 x 10^-5 mol dm^-3)

Calculate the pH of a buffer solution made by mixing together 130 cm³ of 0.2 mol dm^-3 ethanoic acid and 85 cm³ of 0.45 mol dm^-3 sodium ethanoate. Give your answer to 2 decimal places.

(K_a of ethanoic acid = 1.74 x 10^-5 mol dm^-3)

Calculate the new pH of the buffer solution given in part (b) if 1.00 cm³ of 1.00 mol dm^-3 sodium hydroxide is added. Comment on the pH value after the addition of sodium

Buffers are solutions that can resist changes in pH if a small volume of acid or base is added.

i) Write an equation to show the action of the buffer given in part (b) if a small volume of acid was added.

ii) Write an equation to show the action of the buffer given in part (b) if a small volume of base was added.

This question is about pH titration curves.

Figure 1 shows a sketch of a graph pH curve when ammonia, NH₃ (aq), is added to a hydrochloric acid, HCl (aq). 

20.00 cm³ of 0.5 mol dm^-3 HCl (aq) was required to neutralise 38.65 cm³ of NH₃ (aq). 

Figure 1

i) Write two equations to show how hydrochloric acid and ammonia can act as a Brønsted-Lowry acid and base respectively.

[1]

ii) Suggest a value for the pH at point 2 on the curve and explain your answer.

[1]

iii) Calculate the pH of the solution at point 1. Give your answer to 2 decimal places.

[1]

The 'half volume method' can be used to find the K_a of a weak acid as pK_a = pH at the half equivalence point. This method was carried out as follows:

Step 1 15.00 cm³ of a weak acid was added to a conical flask with a few drops of phenolphthalein indicator.

Step 2 This solution was then titrated against NaOH (aq) until the endpoint.

Step 3 The volume of NaOH (aq) required to reach the endpoint was noted.

Step 4 The titration was then repeated without an indicator and adding only half the volume of NaOH (aq) required to reach the endpoint.

Step 5 The pH of the solution was measured using a pH probe.

The student’s graph is shown below in Figure 2.

Figure 2

Using Figure 2, calculate the value for K_a of the acid. Give your answer to 3 significant figures.

The complete neutralisation of sodium carbonate with hydrochloric acid happens in two separate stages. In the first stage, sodium hydrogencarbonate is produced. This then goes on to react further with hydrochloric acid in the second stage.

Figure 3 shows the change in pH during these two stages.

Explain using equations the two reactions in this double endpoint titration.

Figure 3