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Question 1
A chemist plans to make up a buffer solution with a pH of 5.00. The chemist adds solid sodium ethanoate, CH3COONa, to 500 cm3 of 0.250 mol dm-3 ethanoic acid. Ka for ethanoic acid = 1.75 × 10-5 mol dm-3.
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a) | Calculate the mass of sodium ethanoate that the chemist needs to dissolve in the ethanoic acid to prepare this buffer solution. Assume that the volume of the solution doesn't change on dissolving the sodium ethanoate.
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b) | The chemist prepares another CH3COOH/CH3COO- buffer solution containing 0.0300 mol of sodium ethanoate dissolved in 800 cm3 of 7.00 x 10-2 mol dm-3 ethanoic acid. |
Calculate the pH of the solution formed. Give your answer to 2 decimal places.
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c) | A sample of 10.0 cm3 of 2.00 mol dm-3 hydrochloric acid is added to this buffer solution. |
Calculate the pH of the buffer solution after this addition. Give your answer to 2 decimal places.
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Question 2
A buffer solution with a pH of 4.51 is prepared using ethanoic acid, CH3COOH, and sodium ethanoate, CH3COONa. In the buffer solution, the concentration of ethanoate ions is 0.186 mol dm-3.
Ka of ethanoic acid at 298 K = 1.74 x 10-5 mol dm-3.
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a) | Calculate the concentration of ethanoic acid in the buffer solution. Give your answer to 3 significant figures.
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b) | In a different buffer solution, the concentration of ethanoic acid was 0.390 mol dm-3 and the concentration of ethanoate ions was 0.234 mol dm-3. A 5.00 cm3 sample of 1.40 mol dm-3 sodium hydroxide was added to 500 cm3 of this buffer solution. |
Calculate the pH of the buffer solution after the sodium hydroxide was added. Give your answer to two decimal places.
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Question 3
Acid–base titration pH curves can be used to help choose suitable indicators for the following titrations:
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a) | Sketch a pH acid–base titration curve to show how the pH changes during titration A.
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b) | Sketch a pH acid–base titration curve to show how the pH changes during titration B.
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c) | Explain how the choice of indicator is linked to the pH curve.
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Question 4
In a titration, a solution of potassium hydroxide was added gradually from a burette to 20 cm3 of 0.080 mol dm-3 propanoic acid at 298 K. The pH was measured and recorded at regular intervals. The results are shown in the graph below.
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a) | Use the graph to deduce the volume of KOH added at the end-point of the titration.
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b) | Give the expression for the acid dissociation constant, Ka, of propanoic acid, CH3CH2COOH.
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c) | Use the graph to determine the value of Ka for propanoic acid at 298 K. Give your answer to 2 significant figures.
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d) | The table below shows data about some indicators. |
Explain which indicator(s) from the table would be suitable for this titration of propanoic acid with potasium hydroxide.
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e) | A beaker contains 30 cm3 of a buffer solution at pH = 5.0. Two drops of each of the five indicators in the table are added to this solution.
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State the colour of the mixture of indicators in this buffer solution. Assume that the indicators do not react with each other.
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Question 5
L and M are acids. Y and Z are bases. The table below shows the strength of each acid and base. The two acids were titrated separately with the two bases using methyl orange as indicator. The titrations were then repeated using phenolphthalein as indicator. The pH range for methyl orange is 3.1– 4.4.
The pH range for phenolphthalein is 8.3 –10.0.
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a) | Select the correct statement about the indicator(s) that would give a precise end-point for the titration of: |
i) | Acid L with base Y. A both indicators give a precise end-point. B only methyl orange gives a precise end-point C only phenolphthalein gives a precise end-point D neither indicator gives a precise end-point
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ii) | Acid M with base Y. A both indicators give a precise end-point. B only methyl orange gives a precise end-point C only phenolphthalein gives a precise end-point D neither indicator gives a precise end-point
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iii) | Acid L with base Z. A both indicators give a precise end-point. B only methyl orange gives a precise end-point C only phenolphthalein gives a precise end-point D neither indicator gives a precise end-point
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b) | Explain, why the end point of the reaction between acid M and base Z would be difficult to judge accurately using an indicator.
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Question 6
Water dissociates slightly according to the equation: H2O(I) ⇌ H+(aq) + OH-(aq)
The ionic product of water, Kw, is given by the expression Kw = [H+][OH-]
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a) | Explain why the expression for Kw does not include the concentration of water.
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b) | The graph below shows how Kw varies with temperature. |
b) | Explain, with reference to equilibrium position, why the value of Kw increases as the temperature increases.
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c) | Use the graph above to calculate the pH of pure water at body temperature, 37°C. Give your answer to 1 decimal place.
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d) | Calculate the pH of a 0.18 mol dm-3 solution of KOH at 37°C. Give your answer to 1 decimal place.
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Question 7
Methanoic acid, HCOOH is added to butanoic acid, CH3(CH2)2COOH. A reaction takes place to form an equilibrium mixture containing two acid-base pairs. The Ka of methanoic acid is 1.82 x 10-4 mol dm-3 The Ka of butanoic acid is 1.51 x 10-5 mol dm-3
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a) | Complete the equilibrium equation: HCOOH + CH3(CH2)2COOH ⇌ ............................................ + .............................................
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b) | For each of the 3 equilibria shown below, write 'A' or 'B' on the dotted lines below each equation to indicate whether the substance is acting as a Brønsted–Lowry acid (A) or a Brønsted–Lowry base (B).
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c) | A 25.0 cm3 sample of 7.50 x 10-2 mol dm-3 hydrochloric acid was placed in a beaker. Distilled water was added until the pH of the solution was 1.35.
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Calculate the total volume, in cm3, of the solution formed. Give your answer to 3 significant figures.
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d) | Calculate the concentration of hydroxide ions, in mol dm-3, in the hydrochloric acid solution after the addition of distilled water. Give your answer to 3 significant figures. Kw = 1.00 × 10-14 mol2 dm-6
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Question 8
a) | State the meaning of the term weak Brønsted-Lowry acid.
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b) | Boric acid, H3BO3, is a weak Brønsted acid which dissociates into ions in three stages in aqueous solution. The equation for the first dissociation is: H3BO3(aq) ⇌ H+(aq) + H2BO3-(aq) The pKa for this dissociation is 9.24.
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b) | Calculate the pH of a 6.50 x 10-2 mol dm-3 solution of boric acid from the pKa value for the first dissociation. Give your answer to 2 decimal places.
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c) | Explain the two approximations that are made in the calculation in part b).
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d) | Boric acid can undergo futher dissociation. What is the conjugate acid of the HBO32- ion? A H3O+ B H3BO3 C H2BO3- D BO32-
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e) | What is the conjugate base of the HBO32- ion? D BO32- C HBO32- B H3BO3 A OH-
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Question 9
A student adds 50.0 cm3 of 0.250 mol dm-3 methanoic acid, HCOOH, to 50.0 cm3 of 0.0500 mol dm-3 sodium hydroxide. A buffer solution forms. |
a) | Explain what is meant by the term buffer solution.
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b) | Explain why a buffer solution forms.
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c) | Calculate the pH of this buffer solution at 298 K. The Ka of methanoic acid is 1.78 x 10-4 mol dm-3 at 298 K. Give your answer to 2 decimal places.
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d) | The student prepares another buffer solution using the same concentrations of methanoic acid and sodium hydroxide as above. They combine 25.0 cm3 solution of methnanoic acid with 25.0 cm3 solution of sodium hydroxide. |
Suggest whether the pH of the buffer solution would be the same, greater than, or less than the pH calculated in part c). Explain your reasoning.
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