| 1 | 2 | 3 | 4 | 5 | 6 |
| 7 | 8 | 9 | 10 | 11 |
Question 1
A saturated solution of manganese hydroxide, Mn(OH)2, at 298 K has a solubility of 3.6 x 10-5 mol dm-3. |
a) | Write the equilibrium equation, including state symbols, for the changes that happen in a saturated solution of manganese hydroxide in the presence of some solid manganese hydroxide.
|
b) | Calculate the solubility product, Ksp, for manganese hydroxide. State the units of Ksp. Give your answer to 2 significant figures.
|
c) | Water with a manganese concentration above 1.8 x 10-6 mol dm-3 will cause clothes to stain when being washed. The precipitation of Mg(OH)2 by the addition of OH- ions reduces the concentration of Mn2+ in solution. |
Using your answer to part a), calculate the minimum pH at 298 K required to prevent clothes from being stained. Give your answer to 1 decimal place. The ionic product of water, Kw = 1.00 x 10-14 mol2 dm-6 at 298 K.
|
|
Question 2
a) | Calculate the solubility of calcium sulfate, CaSO4 in mol dm-3. Give your answer to 3 significant figures. The solubility product, Ksp, for calcium sulfate is 2.00 x 10-5 mol2 dm-6.
|
b) | Calcium nitrate and potasium sulfate react together as shown by the equation below. Ca(NO3)2(aq) + K2SO4(aq) ➔ CaSO4(s) + 2KNO3(aq) Equal volumes of 2.50 x 10-3 mol dm-3 solutions of calcium nitrate and potassium sulfate are mixed together. |
Predict, with reasoning, whether calcium sufate will precipitate.
|
c) | A saturated solution of magnesium hydroxide, Mg(OH)2, contains the following equilibrium: Mg(OH)2(s) ⇌ Mg2+(aq) + 2OH-(aq) |
Explain what would happen to the solubility of magnesium hydroxide if the following substances are added to the solution:
|
|
Question 3
a) | State what is meant by the term partition coefficient.
|
b) | Ammonia is soluble in both water and organic solvents. An aqueous solution of ammonia is skaken with the immiscible organic solvent trichloromethane. The mixture is left to reach equilibrium. NH3(aq) ⇌ NH3(organic) Samples are taken from each layer and titrated with dilute hydrochloric acid.
|
Calculate the partition coefficient, Kpc of ammonia between trichloromethane and water. Give your answer to 3 significant figures.
|
c) | Using your answer to part b), explain whether ammonia is more soluble in water or in trichloromethane.
|
d) | Butylamine, C4H9NH2, is also soluble in both water and organic solvents. Suggest how the numerical value of Kpc for butylamine between trichloromethane and water would compare to the Kpc value calculated in part b). Explain your answer.
|
|
Question 4
a) | Explain what is meant by the term buffer solution.
|
b) | Carbonic acid, H2CO3, is a weak acid which, together with hydrogencarbonate ions, HCO3-, acts as a buffer to maintain the pH of blood.
|
Explain, in terms of equilibrium, how the carbonic acid–hydrogencarbonate mixture acts as a buffer in the control of blood pH.
|
c) | Healthy human blood needs to be maintained at a pH of 7.40 for the body to function normally. The pKa value for the dissociation of carbonic acid is 6.38.
|
Calculate the ratio of [HCO3-] : [H2CO3] in healthy blood.
|
|
Question 5
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 why a buffer solution forms.
|
b) | 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.
|
c) | 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 b). Explain your reasoning.
|
|
Question 6
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.
|
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.
|
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.
|
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.
|
|
Question 7
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.
|
a) | Calculate the concentration of ethanoic acid in the buffer solution. Give your answer to 3 significant figures.
|
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.
|
|
Question 8
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.
|
a) | Use the graph to deduce the volume of KOH added at the end-point of the titration.
|
b) | Use the graph to deduce the pH of the solution at the half-neutralisation point.
|
c) | Give the expression for the acid dissociation constant, Ka, of propanoic acid, CH3CH2COOH.
|
d) | Use your answers to parts b) and c) to determine the value of Ka for propanoic acid at 298 K. Give your answer to 2 significant figures.
|
|
Question 9
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-]
|
a) | Explain why the expression for Kw does not include the concentration of water.
|
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.
|
c) | Use the graph above to calculate the pH of pure water at body temperature, 37°C. Give your answer to 1 decimal place.
|
d) | Calculate the pH of a 0.18 mol dm-3 solution of KOH at 37°C. Give your answer to 1 decimal place.
|
|
Question 10
a) | State the meaning of the term weak Brønsted-Lowry acid.
|
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.
|
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.
|
c) | Explain the two approximations that are made in the calculation in part b).
|
d) | Boric acid can undergo futher dissociation. What is the conjugate acid of the HBO32- ion? A H3O+ B H3BO3 C H2BO3- D BO32-
|
e) | What is the conjugate base of the HBO32- ion? D BO32- B H3BO3 C HBO32- A OH-
|
|
Question 11
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
|
a) | Complete the equilibrium equation: HCOOH + CH3(CH2)2COOH ⇌ ............................................ + .............................................
|
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).
|
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.
|
Calculate the total volume, in cm3, of the solution formed. Give your answer to 3 significant figures.
|
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
|
|
| 1 | 2 | 3 | 4 | 5 | 6 |
| 7 | 8 | 9 | 10 | 11 |