SK016

Past Year Examination Questions Unit 7

56

Unit 7: Ionic Equilibria

Jun 99

1. (a) What is the pH of a mixture of solutions containing 9.60 mL 0.1 M NaOH solution and

10.00 mL 0.1M HCl solution? <2.69>

(b) The solubility product, Ksp for AgCl is 1.0x1010 M2. Calculate the solubility of AgCl in

i. water <1.0x105

ii. 0.2 M HCl <5.0x1010>

iii. 0.1 M AgNO3 <1.0x109>

2. (a) Define the equivalence point and end point for a titration.

Why the end point of titration must be the same as the equivalence point in a titration?

(b) Calculate the acid ionisation constant, Ka of ethanoic acid , CH3COOH in a mixture of

solution (pH = 5.44) containing 0.1 mol dm3 ethanoic acid solution and 0.5 mol dm3

sodium ethanoate.

What reaction will take place if a small amount of aqueous solution of strong acid is

added to this mixture of solutions? <1.82x105>

(c) At 25C, 2.20% of 0.125 M benzoic acid, C6H5COOH ionized. Write the ionisation

equation for benzoic acid in water. Determine the acid ionisation constant, Ka for benzoic

acid and calculate its pH value. <6.2x105, 2.56>

Jan 00

3. (a) Write the solubility equilibrium equation and calculate the molar solubility for Hg(OH)2

if Ksp = 3.2x1026. <2.0x109>

(b) Write the equation for the base ionisation of 0.5 M aqueous solution of CH3NH2 and

calculate its pH value. [Kb = 3.7x104 M] <12.13>

4. (a) Calculate the pH value for 500 cm3 0.10 M hydrazinium chloride, N2H5Cl solution.

If the above solution was added to 0.20 M 500 cm3 hydrazine, N2H4, what is the new pH

value? [Kb N2H4 = 1.70x107 M, Kw = 1.0x1014 M2] <4.12><7.53>

(b) With the aid of reaction equations, explain what will happen to the solution pH when a

small amount of strong acid and strong base is added to the above mixture of solutions.

Jul 00

5. (a) Compare the solubility of silver chloride salt, AgCl in water and in dilute HCl solution.

Explain.

(b) Does a precipitate form when 1.0x102 mol of Ba(NO3)2 and 2.0x102 mol of NaF are

dissolved in water and made up to 1.0-L solution? [Ksp BaF2 = 1.7x106] <4.0x106>

6. (a) Define the Bronsted-Lowry acid and base.

(b) Explain how a 50-mL HCl solution with pH 1.25 can be prepared from 11.6 M HCl. <0.24>

(c) In an acid-base titration, 10 mL of 0.45 M HCl was added to 40 mL 0.10M NaOH.

Determine the pH of the solution formed. <2.00>

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Past Year Examination Questions Unit 7

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7 (a) Identify the conjugate acid-base pairs in the equation:

i. H2PO

4 + NH3 HPO

2

4 + NH

4

ii CO2

3 + H2O HCO

3 + OH (2M)

(b) Calculate the pH for 0.02 M Ba(OH)2 solution. <12.6>

(c) 0.10 M NaOH solution was titrated with 0.10 M HCN.

i. Write the equation for the reaction between NaOH and HCN.

ii. Explain by using equations of reactions, why the pH of the solution is not 7.0?

iii. Calculate the pH of the solution at the end point. <11.05>

[Ionisation constant of HCN, Ka = 4.0x1010]

Mac 01

8. The table below shows the base ionisation constants, Kb, for some selected compounds.

Compound Kb

C6H5NH2 3.8x1010

N2H4 1.7x106

NH3 1.8x105

NH2OH 1.1x108

i. Arrange the compound in order of increasing base strength.

ii. Give the structure of conjugate acid for each compound and arrange them in order

of increasing acid strength.

9. 20.00 mL 0.15 M NH3 was titrated with 0.20 M HCl till equivalence point.

Explain whether the solution formed is acidic, neutral or basic.

Hence calculate the pH of the solution at equivalence point.

[Kb NH3 = 1.8x105] <5.16>

Aug 01

10. Give the meaning of buffer solution.

Explain how a mixture of acetic acid, CH3COOH and sodium acetate salt, CH3COONa can

function as a buffer solution.

Show how the following correlation can be obtained.

pH = pKa + ]COOHCH[

]COO[CHlog

3

3

11. Determine whether precipitation would occur when 100 mL 1.00x103 M MgCl2 solution was

added to 400 mL 1.50x103 M NaOH solution. [Ksp Mg(OH)2 = 7.10x1012 M3] <2.88x1010>

Aug 02

12. Define pH.

The pH value for a solution containing 0.245 mol HF in 500 mL water is 1.88. Calculate the

ionisation constant for HF. <3.64x104>

13. The OBr ion is the conjugate base for the weak acid HOBr. When 0.10 M NaOBr was

dissolved in water, the pH value was 10.85.

a. Write the equation for the hydrolysis of NaOBr.

b. Calculate the equilibrium constant for the hydrolysis of NaOBr. <5.01x106>

c. Determine the equilibrium constant, Ka for HOBr acid. <2.0x109>

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Past Year Examination Questions Unit 7

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14. (a) State the differences between solubility and solubility product.

(b) The solubility of Ag2SO4 in water at temperature 25C is 0.506 g for every 100-mL

solution. Calculate the solubility product at temperature 25C. <1.70x105 M3>

(c) If an aqueous solution of Na2SO4 was added progressively into an aqueous solution of

0.01 M Ag2SO4, determine the minimum concentration of Na2SO4 needed just to begin

the precipitation of Ag2SO4. <0.0325 M>

Sept 03

15. (a) At temperature 25C, 0.02 M hydrazine, N2H4 solution is 0.69% ionised. Calculate the

i. [OH-] <1.38x104 M>

ii. ionisation constant, Kb <9.59x107 M>

(b) The solubility of PbSO4 in water is 0.038 g L1 . Calculate the

i. solubility product of PbSO4. <1.57x108 M>

ii. solubility of PbSO4 in 0.01 M Pb2+ solution. <1.57x106 M>

16. 1.0 mL of 0.1 M NaOH was added into 1.0 L buffer solution containing 0.15 M sodium

ethanoate and 0.1 M ethanoic acid, CH3COOH. Calculate the change in pH of the solution.

[Ka CH3COOH = 1.75 x 10-5 ] <4.933>

Oct 04

17. Calculate the mass of hydrogen chloride gas, HCl which must be dissolved into 500 mL of a

0.1-M solution of sodium acetate to produce a buffer solution of a pH of 4.74.

[Ka for acetic acid = 1.75x105] <0.93>

18. The solubility product, Ksp of lead iodide, PbI2 is 7.1x109. Calculate the molar solubility of

PbI2 in water and in a solution of 0.10 M Pb(NO3)2. Explain the difference in the molar

solubility values of PbI2 in both cases. <1.21x103, 1.33x104>

19. The solubility of benzoic acid, C6H5COOH in water is 2 g L1. In a food science experiment,

2 g of benzoic acid was added into 500 mL of water as an additive.

i. What are the ionic species present in the solution?

ii. Calculate the pH of the solution. <2.99>

iii. Calculate the amount of benzoate ions, C6H5COO as a percentage weight of the initial

benzoic acid added into the water. [Ka = 6.5x105] <3.12%>

Oct 05

20. The ionisation of phenylacetic acid, C6H5CH2COOH is as follows:

C6H5CH2COOH + H2O H3O+ + C6H5CH2COO Ka = 4.9x105

i. Calculate the concentration of C6H5CH2COO ion in 0.186 M solution of

C6H5CH2COOH. <3.02x103>

ii. What is the pH of this solution? <2.52>

21. The pH of a saturated solution of magnesium hydroxide, Mg(OH)2 is 10.52.

Calculate the molarity of Mg(OH)2 in the solution and its solubility product.

<1.66x104; 1.82x1011>

22. Determine the pH of the solution formed when 15 mL of 0.25 M sodium hydroxide solution,

NaOH, mixes with 25 mL of 0.10 M hydrochloric acid, HCl. <12.49>

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23. A student is asked to prepare a buffer solution at pH 4.6 using 50.00 mL of 0.5 M benzoic

acid, C6H5COOH and sodium benzoate, C6H5COONa.

Calculate the mass of sodium benzoate required to prepare the buffer solution.

[Ka(C6H5COOH) = 6.5x105]

Explain the buffering effect of adding a small amount of NaOH and HCl respectively into the

buffer solution. <9.32>

Oct 06

24. (a) Define acid and base according to Bronsted-Lowry

Identify Bronsted-Lowry acid and base and their conjugates in the equation below.

NH3 + H2PO4- NH4

+ + HPO4

-

(b) Sodium cyanide, NaCN is salt formed when a strong base, NaOH is reacted with a weak

acid HCN.

i) Write the balance equation to show the reaction between NaOH and HCN. Classify

the salt formed.

ii) What would be the expected pH of the NaCN solution ? Explain the answer using

appropriate equation(s).

25. Phenol , C6H5OH is a weak acid used as a general disinfectant and in the manufacture

of plastic. Calculate the pH and the concentration of all species present ( H3O+, C6H5O-,

C6H5OH and OH- ) in a 0.10 M aqueous solution of phenol.

Calculate the percentage of dissociation.

[ Ka = 1.3 x 10-10 ] <5.4, 3.6 x 10-6, 3.6 x 10-3>

Oct 07

26. (a) i. Differentiate between end point and equivalence point in a titration.

ii. Sketch a titration curve for the titration of 25.0 mL of 1.00 M NH3 with 0.25

M HCl. Indicate the pH corresponding to the equivalence point.

(b) Calculate the solubility in g L-1 for calcium fluoride, CaF2.

[Ksp CaF2 = 3.2 x 10-11] <0.016>

27. (a) What is an indicator?

The properties of some common indicators are given in TABLE 1. Choose the best

indicator for detecting the equivalence point of a titration of a week acid and a strong base.

Explain your answer.

TABLE 1

Indicator pKa Effective colour

range pH

Colour in acid

form

Colour in base

form

Methyl orange 4.2 3.1 – 4.4 Red Yellow

Bromothymol blue 7.1 6.0 – 7.8 Yellow Blue

Thymol blue 8.2 7.9 – 9.4 Yellow Blue

Phenolphthalein 9.5 8.3 – 10 Colourless Red

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(b) A 25.00 mL of 0.50 M acetic acid, CH3COOH solution is titrated with 0.50 M sodium

hydroxide, NaOH solution. Calculate the initial pH of the acid solution.

Qualitatively, predict the pH of the solution after the addition of 25.00 mL NaOH solution.

Explain your answer.

[Ka(CH3COOH) = 1.8 x 10-5] <2.52>

2008/2009

28. (a) Calculate the pH of a solution prepared by dissolving 0.26g of sodium hydroxide, NaOH,

in water using a 250 mL volumetric flask. <12.41>

(b) Pyridine, C5H5N, is a weak base. Calculate the percentage ionisation of pyridine if the

initial concentration of pyridine is 5.00 × 10-4 M. <0.17>

An amount of 0.01 mol of concentrated ammonia, NH3, was added to the above pyridine

solution to give a final volume of 500 mL. Calculate the concentrations of ammonium and

pyridinium ion after the addition of the ammonia.

[Kb C5H5N = 1.52 × 10-9, Kb NH3 = 1.8 × 10-5] <6.0 x 10-4 ,1.27 x 10-9>

2009/2010

29. (a) Calculate the pH of a mixture of 50cm3 of ammonia solution of a concentration of

0.10 moldm-3 and 50 cm3 of hydrochloric acid solution of a concentration 0.050

moldm-3

(Kb of ammonia is 1.8x10-5) <9.26>

(b) Choosing suitable examples from the following:

NH3,Cu2+,H2SO4-,NH2

-,H2O

Explain, using a different equation in each case, the meaning of the terms below:

i. Lewis acid

ii. Bronsted-Lowry acid

iii. Conjugate acid-base pair (identify each member of both acid-base

pairs)

30. The solubility products constant ,Ksp for a series of hydroxides are shown below:

Hydroxides Ksp

CuI 1.0x10-12

AgI 8.3x10-17

PbI2 7.1x10-9

BiI3 8.1x10-19

Calculate the molar solubility and list these four compounds in order of decreasing

molar solubility in:

i. Water <1x10-6, 9.1x10-9,1.2x10-3,1.3x10-5>

ii. 0.10 M NaI <1x10-11, 8.3x10-16,7.1x10-7,8.1x10-16>

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31. An amine, RNH2, is a weak base with a dissociation constant, Kb of 1.8 x 10-5. It was

dissolves in water to form a dilute solution

(a) Write the dissociation equation for the base and derive its dissociation constant

expression. [4 marks]

(b) Calculate the pH of a 0.4 M aqueous solution of the amine. <11.43> [6 marks]

32. (a) Sketch the curve for the titration of 30.0 mL of 0.10 M NaOH with 0.10 M HCl. Indicate

the approximate pH at the beginning of the titration and at the equivalence point. What is

the total volume of the solution at the equivalence point? [7 marks]

(b) The solubility product of lead (II) bromide, PbBr2 is 8.0 x 10-5 at room temperature.

Determine the molar solubility of lead (II) bromide, PbBr2 in 0.10 M sodium bromide,

NaBr solution. <8.0 x 10-3> [8 marks]

2011/2012

33. (a) Identify the acid-base conjugate pairs in the following reactions:

i. H2PO4- + NH3 HPO4

2- + NH4+

ii. HClO + CH3NH2 CH3NH3+ + ClO-

(b) The dissociation constant, Ka for 0.10 M hydrofluoric acid,HF solution is 6.77

x 10-4. Calculate the pH of the solution and the ionization percentage of HF. <2.08, 8.23>

(c) The solubility product, Ksp of Ag2SO4 is 1.5 x 10-5. Calculate the solubility of

the salt in distilled water and in 1.0 M H2SO4 solution. <0.0155, 1.94 x 10-3>

[15 marks]

34. (a) A buffer solution was made by dissolving a mixture of sodium acetate,

CH3COONa and acetic acid, CH3COOH in water. Explain how the buffer

maintains the pH of a solution when a small amount of strong acid or a

strong base is added to the buffered solution.

An amount of 0.984 g of CH3COONa was dissolved in 100 ml of 0.1 M

CH3COOH to make an acidic buffer. Calculate the pH of the buffer and

state all the assumptions made.

[Ka CH3COOH = 1.8 X 10-5] <4.82>

[16 marks]

(b) Lactic acid, HL is a monoprotic acid. An aqueous solution of 0.1 M lactic

acid has a pH of 2.34. Calculate the Ka value for this acid. <1.38 x 10-4>

[4 marks]

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2012/2013

35. (a) Define a buffer solution.

(b) Explain how the pH of a buffer solution made from a weak acid and its

conjugate base changes when;

(i) the acid dissociation constant, Ka, of the weak acid increases.

(ii) the acid concentration is decreased relative to the concentration of its

conjugate base. [6 marks]

(c) Calculate the pH of a buffer solution prepared by adding 5.0 g of sodium

acetate, CH3COONa, to 250 mL of 0.20 M acetic acid, CH3COOH. What is

the pH of the buffer solution after addition of 0.01 M sodium hydroxide,

NaOH, solution?

[Ka acetic acid = 1.8 x 10-5] <4.83>

[14 marks]

36. (a) What is meant by end point and equivalence point? [4 marks]

(b) A titration of 25.0 mL solution of 0.15 M HNO3 requires 37.5 mL of NaOH to

reach the equivalence point

(i) Calculate the pH of the solution before the addition of NaOH.

(ii) Calculate the concentration of NaOH solution at equivalence

point.

[4 marks]

(c) A solution contains 0.175 M manganese (II) ions and 0.182 M lead (II) ions.

The Ksp for manganese (II) sulphide is 1.4 x 10-15 and for lead(II) sulphide is

8.0 x 10-28.

(i) Write the solubility product expression for both sulphides.

[4 marks]

(ii) Calculate the maximum concentration of sulphide ions present in the

solution without manganese ions being precipitated.

[3 marks]

2013/2014

37. (a) Define buffer solution. [2 marks]

(b) Explain how the pH of a mixture of aqueous ammonium chloride and

aqueous ammonia remains constant when a small amount of strong base

is added. Give suitable equations in your explanation.

[4 marks]

(c) A buffer solution of pH 4.84 was prepared by dissolving a certain amount

of sodium ethanoate in 1 dm3 of 0.2 M ethanoic acid.

(i) Calculate the mass of sodium ethanoate needed to prepare the

above buffer solution. <20.6>

(ii) Determine the change in pH when 1.0 cm3 of 1.0 M hydrochloric

acid is added to 100 cm3 of the buffer solution.

[Ka for CH3COOH = 1.8 x 10-5 M] <0.04> [9 marks]

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38. In a titration experiment at 25oC, a total of 30 ml of 0.1 M NaOH was added drop

wise into a conical flask containing 25 ml of 0.1 M HCl and a few drops of an

indicator. Show the variation of pH of the solution before the addition , at half

equivalence point, at equivalence point and at the final volume of the titration. Sketch

a graph of pH against the volume of NaOH.

[ Given : KW at 25oC = 1 x 10-14 ] <1.0, 1.5, 7.0, 12.0>

What is acid-base indicator? By referring to TABLE 2 below, suggest the most

suitable indicator to be used in the titration above and give reasons.

TABLE 2

Indicator pH range

Phenolphthalein 8.3 – 10.0

Cresol red 7.0 – 8.8

Bromothymol blue 6.0 – 7.6

Litmus 4.7 – 8.3

[20 marks]

2014/2015

39. A sample of 0.1000 g of NaOH is dissolved in 25 mL of water and titrated stepwise

until 30 mL of 0.100 M HCl is added. Using these data, describe the titration process

from beginning till the end. Sketch a graph showing the titration curve for this

process.

If HCl is replaced by a weak monoprotic acid, HY, sketch the expected titration curve

on the same graph as above. [20 marks]

40. (a) HA is a weak acid with an acid dissociation constant, Ka = 2.95 x 10-8. If the

concentration of the acid is 1.12 M, calculate

(i) pH of the solution. [6 marks]

(ii) percent dissociation. [2 marks]

(b) The solubility product constant, Ksp, of calcium phosphate, Ca3(PO4)2, in pure

water is 1.2 x 10-26 at 25 0C.

(i) Write the expression for the solubility product constant. [2 marks]

(ii) Calculate the molar solubility of calcium phosphate and concentration of

each ion. [5 marks]

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2015/2016

41. (a) Figure 1 shows a titration curve for 20.0 mL of an unknown concentration of

NaOH solution titrated against a standard solution of 1.00 M HCl.

FIGURE 1

(i) Define standard solution, equivalent point and end point.

(ii) What is the pH at equivalent point, Z for this titration?

(iii) Calculate the molarity of the NaOH solution used in the titration.

(iv) Calculate the pH of the solution after the addition of 25.0 mL of HCl.

[8 marks]

(b) Pyridine, C5H5N, is a weak base. This base ionises in water to produce

C5H5NH+ and OH-. Calculate the pH of a 1.00 M pyridine solution if the

dissociation constant, Kb, of pyridine is 1.50 x 10-9.

[7 marks]

42. (a) Milk of magnesia consists of saturated solution of Mg(OH)2 and gelatinous

Mg(OH)2. Determine the molarity og Mg2+ in a saturated solution of Mg(OH)2

and in the saturated solution of Mg(OH)2 containing 0.01 M NaOH.

[ Given : Ksp Mg(OH)2 = 8.9 x 10-12] [6 marks]

(b) A buffer solution is prepared by dissolving 0.125 mol of sodium nitrite,

NaNO2, in 500 mL of 0.25 M nitrous acid, HNO2. What is the pH of the

solution?

The above buffer solution can resist the changes in pH when a small amount of

the strong acid or base is added into it. Explain how this buffer solution

maintains its pH. If an amount of 0.001 mol of NaOH is added to the solution,

determine the change in the pH value of the buffer solution.

[ Given : Ka HNO2= 5.1 x 10-4]

[14 marks]

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2016/2017

43. (a) Methylamine, CH3NH2, is a weak base. This base has the degree of ionization

of 4.7% at 0.16 M concentration. Calculate

(i) the concentrations of OH-, CH3NH3+ and H3O

+

(ii) the pH and

(iii) the base dissociation constant, kb.

(b) Would lead(II)iodide, PbI2, precipitate (Ksp PbI2 = 8.5 X 10-9) be produced

when a 10.0 mL solution of 1.0 M lead(II)nitrate, Pb(NO3)2, and 40.0 mL

solution of 2.0 X 10-3 M sodium iodide, NaI, is mixed?

[15 marks]

44. Name the processes that occur when HNO2 solution is added separately into water,

NaOH solution and NaNO2 solution. Write the chemical equation for each process.

In three separate beakers, 20.0 mL of 0.08 M HNO2 was added to 40.0 mL of water,

40.0 mL of 0.03 M NaOH and 40.0 mL of 0.03 M NaNO2, respectively. Determine

the pH of the solution in each beaker.

[Ka for HNO2 = 4.5 x 10-4]

[20 marks]

2017/2018

45. (a) In a series of acids of the halogen group, hydroiodic acid, HI, is the strongest

acid while hydrofluoric acid, HF, is the weakest.

(i) Explain what meant by strong and weak acids.

(ii) Briefly explain the main factor in determining the strength of the acid.

(iii) Calculate the degree of dissociation, α of 0.20 M HF solution.

[Ka =6.8 x 10-4 at 25 ° C]

[8 marks]

(b) Acetic acids, CH3COOH is a weak acid with Ka = 1.8 x 10-5. It dissociates in

water according to the following equation:

CH3COOH (aq) + H2O (l) CH3COO – (aq) + H3O+ (aq)

(i) Rewrite the above equation and label the conjugate acid and conjugate

base.

(ii) Determine the pH of 0.10 M acetic acid.

(iii) Predict the percentage of dissociation when the acid concentration is

increased.

[7 marks]

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46. (a) Hypochlorous acid, HOCl, is a monoprotic acid. An aqueous solution of 0.028

M HOCl has a pH of 4.5. Calculate the Ka value for this acid.

A buffer solution is prepared by mixing 100 mL of 0.04 M HOCl with 100 mL

of 0.02 M NaOCl. Calculate the pH of the solution. Explain how a mixture of

HOCl and sodium hypochlorite, NaOCl, solution behaves as a buffer solution

when a small amount of strong acid is added.

[12 marks]

(b) The solubility product, Ksp for calcium sulphate, CaSO4, is 2.0 x 10-5 at 25° C.

Calculate the molar solubility of CaSO4 in water in 0.10 M sodium sulphate,

Na2SO4. State common ion effect on the solubility of CaSO4.

[8 marks]

2018/2019

47. (a) A sample of 0.214 g of unknown monoprotic weak acid, HA was dissolved in

25.00 mL of water and titrated with .01 M NaOH. The acid required 27.40 mL

of the base to reach the equivalent point.

(i) Determine the molarity of the acid.

(ii) Calculate the pH of the solution if 40 mL NaOH is added to the acid

solution.

(iii) Explain qualitatively the pH of the solution at equivalent point.

[12 marks]

(b) Magnesium arsenate, Mg3(AsO4)2, is used as an insecticides and it is toxic to

humans. Calculate the solubility of Mg3(AsO4)2 in water at 25°C.

[Given: Ksp Mg3(AsO4)2 = 2.2 x 10-20]

[5 marks]

2019/2020

48. (a) Nitrous acid, HNO2 is used to distinguish between primary, secondary and

tertiary amines. At 25°C, an aqueous solution of 0.215 M HNO2 is 4.0%

dissociated.

(i) Write a balanced equation for the dissociation of HNO2 in water.

(ii) Determine the acid dissociation constant, Ka for HNO2.

(iii) Calculate the pH of the solution.

[6 marks]

(b) Calcium hydroxide, Ca(OH)2 is used to neutralize excess acidity in lakes and

soils. Calculate the pH of 0.30 M Ca(OH)2 solution.

[4 marks]

(c) In a weak acid-strong base titration, 25.00 mL of KOH solution, is titrated

with 21.45 mL of 0.10 M HCN to reach the equivalence point.

(i) Write a balanced chemical reaction equation for the titration.

(ii) Calculate the concentration of KOH solution at the equivalence point.

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(iii) Explain why potassium cyanide, KCN is not a neutral salt using

appropriate equations.

(iv) Suggest a suitable indicator for the titration.

[7 marks]