Sem1 Unit7 Ionic Equilibria

SK017 Unit 7

Past Year Examination Questions

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 1.0x10−10 M2. Calculate the solubility of AgCl in i. water <1.0x10−5 ii. 0.2 M HCl <5.0x10−10> iii. 0.1 M AgNO3 <1.0x10− 9

>

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 dm−3 ethanoic acid solution and 0.5 mol dm−3 sodium ethanoate. What reaction will take place if a small amount of aqueous solution of strong acid is added to this mixture of solutions? <1.82x10− 5

>

c) At 25°C, 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. <6.2x10−5, 2.56> Jan 00 3.

a) Write the solubility equilibrium equation and calculate the molar solubility for Hg(OH) 2 if Ksp = 3.2x10−26. <2.0x10−9>

b) Write the equation for the base ionisation of 0.5 M aqueous solution of CH3NH2 and calculate its pH value. 4.

[Kb = 3.7x10−4 M]

<12.13>

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 = 1.70x10−7 M, Kw = 1.0x10−14 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 1.0x10−2 mol of Ba(NO3)2 and 2.0x10−2 mol of NaF are dissolved in water and made up to 1.0-L solution? [Ksp BaF2 = 1.7x10−6] <4.0x10−6>

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.

33

<2.00>

SK017 Unit 7

Past Year Examination Questions

7

a) Identify the conjugate acid-base pairs in the equation: i. H2PO −4 + NH3 HPO 24− + NH +4 ii

2−

−

CO 3 + H2O → HCO 3 + OH−

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

(2M) <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> −10 [Ionisation constant of HCN, Ka = 4.0x10 ] Mac 01 8.

The table below shows the base ionisation constants, Kb, for some selected compounds. Compoun Kb d C6H5NH2 3.8x10−10 N2H4 1.7x10−6 NH3 1.8x10−5 NH2OH 1.1x10−8 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 = 1.8x10−5]

<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 + log

11.

[CH 3 COO − ] [CH 3 COOH]

Determine whether precipitation would occur when 100 mL 1.00x10−3 M MgCl2 solution was added to 400 mL 1.50x10−3 M NaOH solution. [Ksp Mg(OH)2 = 7.10x10−12 M3] <2.88x10−10>

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. <3.64x10−4>

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. 34

SK017 Unit 7

Past Year Examination Questions

a. Write the equation for the hydrolysis of NaOBr. b. Calculate the equilibrium constant for the hydrolysis of NaOBr. c. Determine the equilibrium constant, Ka for HOBr acid. 14.

<5.01x10−6> <2.0x10−9>

(a) State the differences between solubility and solubility product. (b) The solubility of Ag2SO4 in water at temperature 25°C is 0.506 g for every 100-mL solution. Calculate the solubility product at temperature 25°C. <1.70x10−5 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 25°C, 0.02 M hydrazine, N2H4 solution is 0.69% ionised. Calculate the i. [OH-] <1.38x10− 4

M>

ii. ionisation constant, Kb

<9.59x10−7 M>

b) The solubility of PbSO4 in water is 0.038 g L−1 . Calculate the i. solubility product of PbSO4. ii. solubility of PbSO4 in 0.01 M Pb2+ solution. 16

<1.57x10−8 M> <1.57x10−6 M>

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 = 1.75x10−5] <0.93>

18.

The solubility product, Ksp of lead iodide, PbI2 is 7.1x10−9. 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. <1.21x10−3, 1.33x10−4>

19.

The solubility of benzoic acid, C6H5COOH in water is 2 g L−1. 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 = 6.5x10−5] <3.12%>

Oct 05 20.

The ionisation of phenylacetic acid, C6H5CH2COOH is as follows: C6H5CH2COOH + H2O H3O+ + C6H5CH2COO− Ka = 4.9x10−5 i. Calculate the concentration of C6H5CH2COO− ion in 0.186 M solution of C6H5CH2COOH. <3.02x10−3> 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. 35

SK017 Unit 7

Past Year Examination Questions

Calculate the molarity of Mg(OH)2 in the solution and its solubility product. <1.66x10−4; 1.82x10−11> 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>

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) = 6.5x10−5] Explain the buffering effect of adding a small amount of NaOH and HCl respectively into the buffer solution.

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 ]

Oct 07 26.

(a) i.

ii.

Differentiate between end point and equivalence point in a titration. 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] 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 36

Colour in acid

Colour in base

SK017 Unit 7

Past Year Examination Questions

Methyl orange Bromothymol blue Thymol blue Phenolphthalein

4.2 7.1 8.2 9.5

range pH 3.1 – 4.4 6.0 – 7.8 7.9 – 9.4 8.3 – 10

form Red Yellow Yellow Colourless

form Yellow Blue Blue Red

(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] Oct 08 28.

(a) Calculate the pH of a solution prepared by dissolving 0.26 g of sodium hydroxide, NaOH, in water using a 250 mL volumetric flask.

(b) Calculate the solubility of magnesium phosphate, Mg3(PO4)2, in g L-1. [Ksp Mg3(PO4)2 = 5.2 x 10-24, Mg3(PO4)2 = 262 g mol-1] 29.

Pyridine, C5H5N, is a weak base. Calculate the percentage ionisation of pyridine if the initial concentration of pyridine is 5.00 x 10-4 M. 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 x 10-9, Kb NH3 = 1.8 x 10-5]

37