nss chemistry part 13 industrial chemistry _i

8/12/2019 NSS Chemistry Part 13 Industrial Chemistry _I

1/36


2/36

2

2. [HKAL 1990 I Q2d]

Consider the reaction A(l) + B(l) C(l) + D(l) which is first order with respect to A, but second order with

respect to B.

a. Give a rate expression for this reaction.

b. Draw and label suitable straight line graphs that would allow the rate constant to be determined from a

series of experiments in which:

(i) [A] is kept constant, but [B] is altered;

(ii) [B] is kept constant, but [A] is altered.

(4 marks)

3. [HKAL 1990 II Q1c]

Give, with explanations, two factors that would increase the rate of a reaction.

(4 marks)


3/36

3

4. [HKAL 1992 I Q2a]

For the first order reaction

Ak

products,

the integrated form of the rate equation is [A] = [A]oe-kt, where [A] and [A]oare the concentrations of A at

time = t and time = 0 respectively.

(i) Starting from this equation, derive the relationship between the half-life t1/2

of the reaction and the

rate constant k.

(ii) Without using the half-life method, show how you would determine the rate constant k from a set of

experimental measurements of concentration at various times.

(3 marks)


4/36

4

5. [HKAL 1992 II Q2a]

In acidic solution, chlorate(V) ions, ClO3-, slowly oxidize chloride ions to chlorine. The following kinetic data

are obtained at 25oC.

[ClO3-]

/mol dm-3

[Cl-]

/mol dm-3

[H+]

/mol dm-3

Initial rate

/mol dm-3s-1

0.08 0.15 0.20 1.0 10-5

0.08 0.15 0.40 4.0 10-5

0.16 0.15 0.40 8.0 10-5

0.08 0.30 0.20 2.0 10-5

(i) Write the balanced equation for this reaction.

(ii) Determine the order of the reaction with respect to each reactant.

(iii) Determine the rate constant at this temperature.

(6 marks)


5/36

5

6. [HKAL 1993 II Q2a, b]

a. Discuss, in terms of the Arrhenius equation, the effect of temperature on the rate of a reaction.

(2 marks)

b. Given the following data for the reaction at 298K

2C + 3D + E P + 2Q

Experiment [C]

/mol dm-3

[D]

/mol dm-3

[E]

/mol dm-3

Initial rate

/mol dm-3s-1

1 0.10 0.10 0.10 3.010-3

2 0.20 0.10 0.10 2.410-2

3 0.10 0.20 0.10 3.010-3

4 0.10 0.10 0.30 2.710-2

(i) Deduce the rate law of the above reaction.

(ii) Calculate the rate constant.

(6 marks)


6/36

6

7. [HKAL 1994 I]

In an experiment to determine the rate constant at 298 K for the decomposition of sodium thiosulphate by a

large excess of dilute hydrochloric acid, the time, t, taken for a certain amount of sulphur to appear was

measured. Under these conditions: Rate = k'[S2O32-].

(i) Write a balanced equation for the reaction between hydrochloric acid and sodium thiosulphate.

(ii) Describe how you carry out this experiment, indicating the measurements you would take.

(iii) What quantities would you plot on a graph for the determination of the rate constant k' at 298K?

(6 marks)

(i) S2O32-

(aq) + 2H+(aq) H2O(l) + SO2(g) + S(s) (1)

(ii) - Make S2O32-

solutions of known concentration and place equal volume of these solutions in

conical flasks. (1)

- The conical flask should be placed on top of paper marked with a black X. (1)

- Measure the temperature of the solutions (0.5)

- Add a fixed volume of acid to the S2O32-

solution, start a stop-clock (0.5)

- Note the time, t, when a black X marked on a paper becomes invisible when looking

through the flask. (1)

(iii) 1/t against [S2O32-

] (1)


7/36

7

8. [HKAL 1994 II Q1c]

The table below lists the concentration of the reactant C as a function of time at 298K for the following

reaction.

C D

Time/s 0 60 120 180 240 300

[C]/ 10-2mol dm-3 20.0 16.1 13.5 11.6 10.2 9.1

(i).# Plot two graphs to show that the data fit a second order reaction better than a first order reaction.

(ii)# Determine the rate constant of the reaction at 298 K.

(iii) The rate constant of the above reaction is found to be doubled when the temperature is raised from

298 K to 306 K. Determine the activation energy of the reaction.

[Gas constant R = 8.31 JK-1mol-1]

(7 marks)


8/36

8

9. [HKAL 1995 I Q1d]

The iodination of propanone is catalysed by hydrogen ions. The overall equation is:

CH3COCH3(aq) + I2(aq) CH3COCH2I(aq) + HI(aq)

Using four mixtures B, C, D and E, the progress of the reaction was followed by colorimetry measurement. The

results are tabulated below.

Mixture

Composition by volume of mixture / cm3 Initial rate

/ mol dm-3

s-1

propanone water 1.00M HCl 0.05 M I2in KI

B 10.0 60.0 10.0 20.0 4.96 10-6

C 10.0 50.0 10.0 30.0 5.04 10-6

D 5.0 65.0 10.0 20.0 2.45 10-6

E 10.0 65.0 5.0 20.0 2.47 10-6

(i) Determine the effects of the changes in concentration of each of the reactants (iodine and propanone)

and the catalyst (hydrochloric acid) on the reaction rate. Write an expression for the reaction.

(ii) For mixture B, calculate the rate constant for the reaction at the temperature of the experiment.

(Density of CH3COCH3= 0.789 g cm-3

)

(4 marks)

(i)

(ii)


9/36

9

10. [HKAL 1995 II Q2b]

In an aqueous solution, the decomposition of hydrogen peroxide in the presence of manganese(IV) oxide can be

represented by the following equation:

)()(2)(2 22222 gOlOHaqOH

MnO+

(i) For a given amount of manganese(IV) oxide, outline how you would use a chemical method to

determine the concentration of hydrogen peroxide, at different times, in the reaction mixture.

(Hint: redox titration between acidified potassium permanganate and hydrogen peroxide)

(ii) How would you use the results obtained in (i) to show graphically that the decomposition is first order

with respect to H2O2and to find the rate constant of the decomposition?

(6 marks)


10/36

10

11. [HKAL 1997 II Q3]

The reaction of iodine with propanone in acidic solutions can be represented by the following equation:

CH3COCH3(aq) + I2(aq) CH3COCH2I(aq) + H+(aq) + I

-(aq)

a. (i) The progress of the reaction can be monitored by a titrimetric method. Outline the experimental

procedure.

(ii) State how the initial rate of the reaction can be determined from the titrimetric results.

(iii) Suggest another method to monitor the progress of the reaction.

(6 marks)

b. The following initial rates and initial concentrations were obtained in an experiment at 298K:

Initial rate

/ mol dm-3

s-1

Initial concentration / mol dm-3

[I2(aq)] [CH3COCH3(aq)] [H+(aq)]

3.5 10-5 2.5 10-4 2.0 10-1 5.0 10-3

3.5 10-5

1.5 10-4

2.0 10-1

5.0 10-3

1.4 10-4

2.5 10-4

4.0 10-1

1.0 10-2

7.0 10-5

2.5 10-4

4.0 10-1

5.0 10-3

(i) Deduce the rate equation for the reaction.

(ii) Calculate the rate constant for the reaction at 298K.

(5 marks)

c. Suppose that the reaction takes place in a buffer solution of pH 4. On the basis of your results in (b). deduce the

half-life of the reaction at 298K.

(Note: Buffer solution is used to keep the pH of a solution to be constant.)

(4 marks)

d. For a given set of initial concentrations, the initial rate doubles when the temperature is increased from 298K to

308K. Calculate the activation energy of the reaction.

(2 marks)


11/36

11


12/36

12

12. [HKAL 1998 II Q2b]

Potassium-40 is radioactive and decays to give the stable isotope, argon-40. The half-life of the decay is 1.27

109years.

(i) In a rock sample, the ratio of40

K to40

Ar is 1 to 9. Estimate the age of the rock sample.

(ii) The above method of estimation is bases on several assumptions. One of the assumptions is that all40

Ar

present in the rock sample is derived from the decay of40

K. Give one other assumption relating to40

Ar.


13/36

13

13. [HKAL 1999 II Q3a]

Consider the following data for the reaction:

A + B products


Initial rate

/ mol dm-3

s-1

[A] [B]

4.0 10

-2

4.0 10

-2

6.4 10

-5

8.0 10

-2 4.0 10

-2 12.8 10

-5

4.0 10-2

8.0 10-2

6.4 10-5

For this reaction,

(i) deduce its rate equation,

(ii) calculate the rate constant, and

(iii)# sketch a possible energy profile.

(6 marks)


14/36

14

14. [HKAL 2001 II Q3a]

At an elevated temperature and in the presence of argon, iodine atoms combine to form iodine molecules:

2I(g) I2(g)

The table blow lists some data about the reaction:


Initial rate

dtgId )]([ 2 / mol dm-3s-1[I(g)] / mol dm-3 [Ar (g)] / mol dm-3

1.0 10-5

1.0 10-3

8.70 10-4

2.0 10-5

1.0 10-3

3.48 10-3

1.0 10-5

5.0 10-4

4.35 10-4

For this reaction,

(i) deduce the rate equation,

(ii) calculate the rate constant,

(iii) suggest a possible reaction mechanism consistent with the rate equation#, state the role of argon, and

sketch the energy profile.

(8 marks)

(iii)


15/36

15

15. [HKAL 2002 I Q2b]

An experiment was carried out to study the acid-catalyzed bromination of propanone at 298K.

CH3COCH3(aq) + Br2(aq) +

H CH3COCH2Br(aq) + HBr(aq)


Initial rate /

mol dm

-3

s

-1

[CH3COCH3(aq)] [Br2(aq)] [H

+

(aq)]

0.30 0.050 0.050 5.7 10-5

0.30 0.100 0.050 5.7 10-5

0.30 0.050 0.100 1.2 10-4

0.40 0.050 0.200 3.1 10-4

0.40 0.050 0.050 7.6 10-5

(i) Deduce the rate equation for the reaction.

(ii) Calculate the rate constant for the reaction at 298 K.

(5 marks)

(i)

(ii)


16/36

16

16. [HKAL 2003 II Q3a]

An experiment was carried out at 298K to study the acid-catalyzed hydrolysis of sucrose:

C12H22O11(aq) + H2O(l) HCl

C6H12O6(aq) + C6H12O6(aq)

Sucrose glucose fructose

The table below lists the concentration of sucrose, in arbitrary units, at different times.

Time / minutes 0 60 120 180 240[C12H22O11(aq)] 100.0 81.3 66.3 54.0 44.1

(i) Suggest a method that can be used to monitor the progress of the hydrolysis.

(ii) By plotting a suitable graph, show that the hydrolysis is first order with respect to sucrose. That is, the

rate equation of the hydrolysis can be represented by:

rate of reaction = k [C12H22O11(aq)]

(iii) Using your graph in (ii), determine k at 298K.

(iv) Suggest how you would show experimentally that the rate of the hydrolysis is also first order with

respect to hydrochloric acid.

(8 marks)


17/36

17

17. [HKAL 2004 II Q3a]

a. The decomposition of dinitrogen pentoxide in tetrachloromethane can be represent by the following equation:

2N2O5(in CCl4) 4NO2(in CCl4) + O2(g)

(i) Suggest an experimental method that can be used to follow the progress of the decomposition, and state the

underlying principle of the method.

(ii)# The table below lists the results obtained in an experiment to study the kinetics of the decomposition of

dinitrogen pentoxide at 318 K.

By plotting a suitable graph, determine the rate equation for the decomposition, and hence the rate constant at 318 K.

(iii) The rate

constant for the decomposition at 332 K was found to be 5 times that at 318 K. Calculate the activation

energy for the decomposition.

(8 marks)

time / minutes 0 20 40 60 80

[N2O5] / 102

mol dm3

10.00 5.49 3.02 1.65 0.91


18/36

18

18. [HKAL 2005 II Q4a]

In an experiment to study the kinetics of the reaction:

2I-(aq) + S2O8

2-(aq) I2(aq) + 2 SO4

2-(aq)

the time (t) for the formation of a very small but fixed amount of I2from different mixtures of KI(aq) and

K2S2O8(aq) was measured. The results of three runs of the experiment are listed below:

(i) Describe and explain how time (t) can be found experimentally.

(ii) Deduce from the above information the rate equation for the reaction of I-(aq) with S2O8

2-(aq).

(iii) A fourth run of the experiment was conducted using a mixture with [I-(aq)]oand [S2O8

2-(aq)]oof 0.100 mol

dm-3

and 0.060 mol dm-3

respectively. Calculate the time (t) required for the formation of the same amount

of I2.

(8 marks)

RunInitial concentration / mol dm

-3

t/s[I

-(aq)]o [S2O8

2-(aq)]o

1 0.080 0.080 22.0

2 0.080 0.040 44.0

3 0.040 0.040 88.0


19/36

19

19. [HKAL 2007 II Q3a]Consider the reaction below:

Br2(aq) + HCO2H(aq) 2 Br

(aq) + 2 H+

(aq) + CO2(g)

(i) Suggest an experimental method to follow the change in concentration of Br2(aq) in the reaction mixture.

Give a reason for your suggestion.

The table below lists the experimental data obtained at a certain temperature:

Run

Volume used / cm3 Initial rate for the

disappearance of

Br2(aq) / mol dm3s

1

0.010 M Br2(aq) 0.20 M HCO2H(aq) H2O( )

1 2.0 10.0 8.0 1.2 105

2 4.0 10.0 6.0 2.4 105

3 8.0 10.0 2.0 4.8 105

(ii) Suggest how the initial rate for the disappearance of Br2(aq) can be found.

(iii) Why is it necessary to keep the concentration of HCO2H(aq) much higher than that of Br2(aq)?

(iv) Deduce the order of the reaction with respect to Br2(aq).

(v) Suggest how the order of the reaction with respect to HCO2H(aq) can be determined.

(9 marks)


20/36

20

20. [HKAL 2006 II Q3a]

An experiment is conducted, at 293 K, to study the kinetics of the decomposition of H2O2(aq) in the presence

of peroxidase, an enzyme.

2H2O2(aq) peroxidase

H2O(l) + O2(g)

(i) Outline a chemical method to follow the change in concentration of H2O2(aq) in the reaction mixture at

different times (t).

(ii)# The table below lists the results obtained:

t/ minutes 0 5 10 20 30 50[H2O2(aq)]

(arbitrary units) 46.1 37.1 29.8 19.6 12.3 5.0

By plotting a suitable graph, show that the decomposition is first order with respect to H2O2(aq). Hence,

calculate the rate constant at 293 K.

(iii) Is it possible to determine experimentally the rate constant for the above decomposition at 353 K? Explain.

(9 marks)


21/36

21

21. [HKAL 2008 II 2a]

The gaseous reaction below takes places at 750oC in a closed container with a fixed volume.

(g)NO(g)2H2NO(g)g)(H2 222 ++

(i) Suggest an experimental method that can be used to follow the concentration of (g)N2 in the

reaction mixture. Briefly explain the principle of your suggested method.

(ii) The table below lists three sets of experimental data of the reaction at 750oC:

Experiment

Initial concentration / mol dm3 Initial rate for the

formation of g)(N 2

/ mol-1-3

sdm (g)H2 NO(g)

1 0.010 0.0250 0.500

2 0.005 0.0250 0.250

3 0.010 0.0125 0.125

Deduce the rate equation for the reaction, and calculate its rate constant at 750oC.

(iii) Is the reaction an elementary reaction? Explain.

(6 marks)


22/36

22. [HKAL 2009 II 2a]

An experiment was devised to stu

Four runs of the experiment were

prepared according to the table bel

Run0.02 M I2(aq)

1 50.0

2 50.0

3 50.0

4 50.0

In each run, 10.0 cm3of the reacti

NaHCO3(aq). When effervescence

using starch solution as indicator.

(i) What is the purpose of addin

(ii) Deduce the reaction order wi

(iii) Deduce, by plotting another

(iv) Suggest how the reaction ord

22

y the kinetics of the following acid-catalysed re

onducted at the same temperature, and the reacti

ow:

Volume used /cm3

1.0 M

CH3COCH3(aq)H2O()

5.0 20.0

10.0 15.0

15.0 10.0

20.0 5.0

n mixture was withdrawn at regular time interva

subsided, the resulting mixture was titrated agai

The graph below shows the plot of the titre aga

the reaction mixture to excess NaHCO3(aq) bef

h respect to Iodine.

raph, the reaction order with respect to propanon

r with respect to H+(aq) can be determined.

ction:

on mixtures used were

1.0 M H2SO4(aq)

25.0

25.0

25.0

25.0

ls and added to excess

nst standard Na2S2O3(aq)

inst time for each run.

re each titration?

e.

(11 marks)


23/36

23


24/36

24

Activation Energies

23. [HKAL 1989 I Q2c]

The Maxwell-Boltzmann distribution of molecular speeds of a gas at 1250 K is shown in the diagram below:

speeds/ ms-1

percentage

of

molecules

(i) Sketch on the diagram a graph of the probable distribution of molecular speeds of the same gas at 298K.

(ii) Explain why there is a distribution of speeds.

(4 marks)

(i)

(ii)


25/36

25

24. [HKAL 1991 II Q3a]

For the reaction 2XY(g) X2(g) + Y2(g),

the rate constant is 3.9110-4mol-1dm3s-1at 370oC and 4.05 10-2mol-1dm3s-1at 470oC.

Generally the rate constant of a reaction is related to the temperature by k = A exp(-Ea/RT).

Calculate

(i) the activation energy,

(ii) the rate constant at 450oC.

[Gas constant, R = 8.314 JK-1mol-1]

(6 marks)

25. [HKAL 1992 I Q2e]

The diagram below gives the Maxwell-Boltzmann energy distribution of a system of molecules at two

temperatures:

EaEnergy

Relative no. of

molecules with a

particular energy

T1

T2

where T2 > T1 and Ea is the activation

energy of a reaction What do the shaded areas of the curves represent and why are they different at different temperatures?

(2 marks)

The shaded area under either one of these curves corresponds to the amount of molecules that collide with

kinetic energy which is greater than the activation energy (Ea).

As temperature increases, the average kinetic energy of the molecules will increase, therefore, the amount of

molecules with kinetic energy which is greater than the activation energy (Ea) also increase.


26/36

26

26. [HKAL 1996 I Q1b]

Carbon-14,14

C, is radioactive, emitting particles.

(i)# Write an equation for the decay of14

C.

(ii) A charcoal sample from the ruins of an ancient settlement was found to have a14

C/12

C ratio 0.60 times that

found in living organisms.

(1) Explain why the14

C/12

C ratio in the charcoal sample is smaller than that in living organisms.

(2) Given that the half-life for the decay of14

C is 5730 years, calculate the age of the charcoal

sample.

Note: The integrated form of the rate expression for radioactive decay can be represented by the

following equation:

log.

10

12

0301N

N tto

t

=

where Nois the initial number of radioactive nuclei;

Ntis the number of radioactive nuclei at time t;

t 12 is the half-life for the decay.

(iii) All radioactive decay has zero activation energy. Comment on the effect of temperature upon the rate of

decay.

(5 marks)

(i)

(ii) (1)

(2)

(iii)


27/36

27

27. [HKAL 1998 II Q3a]

The table below lists the rate constants, k, at different temperatures, T, for the first order decomposition of a

dicarboxylic acid, CO(CH2CO2H)2, in aqueous solution:

CO(CH2CO2H)2 (aq)CH3COCH3(aq) + 2CO2(g)

T /K 273 293 313 333 353

k /s

-1

2.46 10

-5

4.75 10

-4

5.76 10

-3

5.48 10

-2

?(i) Determine the activation energy for the reaction by plotting an appropriate graph.

(ii) Estimate the rate constant of the reaction at 353 K and hence calculate the half-life of the reaction at

the same temperature.

(iii) Suggest a method to monitor the progress of the reaction.

(9 marks)


28/36

28

28. [HKAL 1998 II Q3b]

The exothermic reaction

E(g) E(g) (1)

is a single stage reaction.

(i) Sketch curves to show the distribution of molecular kinetic energy of the reactant, E(g), at two

different temperatures.

(ii) With reference to your answer in (i), explain why the rate of reaction (1) increases with temperature.

(iii) In the presence of a catalyst, C, reaction (1) will proceed at a faster rate via the following mechanism:

E(g) + C(g) EC(g)

EC(g) C(g) + E(g) (EC is the reaction intermediate.)

Sketch labeled energy profiles for the conversion of E(g) to E(g), with and without the catalyst.

Explain why reaction (1) proceeds faster in the presence of the catalyst.

(8 marks)

(i)


29/36

29

29. [HKAL 2000 I Q7b]

Without giving any experimental detail, outline what measurements have to be taken in order to determine

the activation energy of a reaction.

(3 marks)

Keeping the (initial) concentrations of the reagents unchanged, carry out the experiment at different

temperatures and determine the corresponding (initial) rate.

(ii)

(iii)


30/36

30

30. [HKAL 2002 II Q3a]

An experiment was conducted to study the hydrolysis of benzenediazonium chloride at 298K and 1 atm.

C6H5N2+Cl

-(aq) + H2O(l) C6H5OH(aq) + N2(g) + HCl(aq)

The progress of the hydrolysis was followed by measuring the volumes of N2(g), v, liberated at different time, t.

The table below lists the experimental results.

t / s 0 5 10 15 25 3600 5400

v / cm3 0 33 56 73 92 110 110

a. By plotting a suitable graph, show that the hydrolysis is first order with respect to benzenediazonium

chloride.

b. Using your graph in (a), determine the rate constant of the hydrolysis at 298K.

c. Suggest how the activation energy of the hydrolysis can be determined.

(You are not required to give the experimental details.)

(8 marks)


31/36

31


32/36

32

31. [HKAL 2004 I Q3a]

a Consider the following system, which comprises two single step reactions.

(k1and k1are the rate constants.)

(i) Write the respective rate equations for the forward and backward reactions.

(ii) Sketch a labeled energy profile for the forward reaction.

(iii) Predict with explanation, whether k1or k1will increase to a greater extent when the temperature of the

system is increased.

(4 marks)


33/36

32. [HKAL 2008 I 3a(iii)]

Which one of the following gra

two different temperatures T1an

33. [HKAL 2010 I 1b]

A student made the following rem

The rate of an elementary gaseou

the reactant molecules increases w

Is the explanation provided by the

Elaborate your answer.

Ans: C

33

hs represents the distribution of molecular spe

T2, where T2> T1?

ark:

s reaction increases with temperature because t

ith temperature.

student regarding the increase in reaction rate ap

ds of one mole of a gas at

e average kinetic energy of

propriate ?

(3 marks)


34/36

34

Catalysts

34. [HKAL 1991 I Q2b]

The energy profiles of the reaction

A(g) + B(g) C(g)

under two different catalysts X and Y are represented below.

Reaction coordinate

Energy

System with catalyst X

A, B

C

Reaction coordinate

Energy

System with catalyst Y

A, B

C

a. What is the effect of increasing temperature on the equilibrium of each system?

b. What is the effect of decreasing pressure on the equilibrium of each system?

c. Compare the effect of increasing temperature on the rate of reaction in the two systems.

d. Why could the use of a different catalyst change the order of the reaction?

(7 marks)


35/36

35

35. [HKAL 1994 Essay]

Write an essay to account for the fact that at 750K and 200 atmosphere pressure, the energetically favourable

formation of ammonia

N2(g) + 3H2(g) 2NH3(g)

proceeds very slowly, but in the presence of an iron surface it proceeds at a much faster rate.

(20 marks)


36/36

36. [HKAL 2002 I Q7b]

Devise an experiment, using chemicals and apparatus commonly available in a school laboratory, to show

that the reaction of peroxodisulphate(VI) ions with iodide ions can be catalyzed by iron(III) ions.

S2O82-

(aq) + 2I-(aq) 2SO4

2-(aq) + I2(aq)

(5 marks)

nss chemistry part 13 industrial chemistry _i

Documents