NATIONAL UNIVERSITY OF SINGAPORE
CM1502 LABORATORY REPORT
EXPERIMENT EXPERIMENT 2: Rate of An Iodine-Clock Reaction
Name: Chai Mei Xian Audrey Matriculation number: U075391A Laboratory Group: 3A
Objectives
1) To determine the values of the rate constant (k) and the orders of reaction (m and n) with respect to each of the reactants for the following reaction, using the method of initial rates. 2 I- (aq) + S2O82- (aq) → I2 (l) + 2 SO42- (aq) Hence, to determine the rate law as given by the following expression, R = k [I-] n [S2O82-] m where R is the rate of reaction. 2) To determine the activation energy E , for the following reaction from the temperature dependence of the rate constant of the reaction. 2 MnO4- (aq) + 5C2O42- (aq) + 16 H+ → 2 Mn2+ (aq) + 10 CO2 (g) + H2O (l) ₐ
Results
Determination of Reaction Orders and Rate Constant Using Methods of Initial Rates
[S2O82-] = 0.100 M [I-] = 0.200 M [S2O32-] = 0.00330 M Temperature = 30.0 oC (kept constant for all runs) Solution
Vol.S 2O82-
Vol. I/ mL
Vol. H2O / mL
Vol. Starch /mL
Vol. S2O32/ mL
Time/s
/ mL
1 2
10.00 10.00
10.00 8.00
0.00 2.00
1.00 1.00
5.00 5.00
19 26
3 4
10.00 10.00
6.00 5.00
4.00 5.00
1.00 1.00
5.00 5.00
39 48
5
10.00
3.00
7.00
1.00
5.00
86
6 7
10.00 8.00
10.00 10.00
0.00 2.00
1.00 1.00
5.00 5.00
21 27
8 9
6.00 5.00
10.00 10.00
4.00 5.00
1.00 1.00
5.00 5.00
38 45
10
3.00
10.00
7.00
1.00
5.00
82
Determination of the Temperature Effect on a Chemical Reaction.
[(COOH)2] = 0.300M [KMnO4] = 0.100M Temperature/ oC
Time/s
10.0
1144
20.0
322
30.0 45.0
86 16
60.0
4
Calculations
1)
Calculation of [I -] and [S2O82-] for each reacting solution
Initial [S2O82-]
= 0.100 mol/L
Initial [I-]
= 0.200 mol/L
Total volume of each solution = 0.026L η S2O82- = vol. of S2O82- used (in mL)/1000 x initial [S2O82-] η I-
= vol. of I- used (in mL)/1000 x initial [I-]
[S2O82-] after dilution = η S2O82- /total vol. of each solution [I-] after dilution
= η I- /total vol. of each solution (where η is the number of moles.)
Solution
Vol. 0f S2O82-/ mL
Vol of I-/ mL
[I-]/M
[S2O82-]/M
1
10.00
10.00
0.0769
0.0385
2
10.00
8.00
0.0615
0.0385
3 4
10.00 10.00
6.00 5.00
0.0462 0.0385
0.0385 0.0385
5 6
10.00 10.00
3.00 10.00
0.0231 0.0769
0.0385 0.0385
7
8.00
10.00
0.0769
0.0308
8 9
6.00 5.00
10.00 10.00
0.0769 0.0769
0.0231 0.0192
10
3.00
10.00
0.0769
0.0115
Calculation of the no. of moles of I 2 that reacted with the S 2O32- per litre per second for each reacting solution Vol. of S2O32- used in each reacting solution = 0.005L [S2O32-] = 0.00330 mol/L
2)
∴ η S2O32- = 0.005 x 0.0033
= 1.65 x 10-5 mol Based on the reaction between I2 and S2O32-, I2 + 2 S2O32- → 2 I- + S4O62η I2 reacted = 0.5 x η S2O32= 0.5 x 1.65 x 10-5 = 8.25 x 10-6 mol Total vol. of each reacting solution = 0.026L η I2 that reacted with S2O32- per litre per second for each reacting solution = (η I2 reacted/ Total vol. of each reacting solution)/ Time taken for solution to turn blue = (8.25 x 10-6 mol/ 0.026L) / Time taken for solution to turn blue = (3.17 x 10-4) / Time taken for solution to turn blue Rate of Iodide-Persulfate Reaction 2I- + S2O82-
I2 + 2S2O32-
I2 + 2SO422I- + S4O62-
∴ Rate of iodide-persulfate reaction = rate of I2 produced
= η I2 that reacted with S2O32- per litre per second (since all the iodine produced per litre per second reacted with S2O32- before excess iodine produced turned solution blue.) = (3.17 x 10-4) / Time taken for solution to turn blue
η I2 that reacted with S2O32per litre per second /mol
Rate of iodidepersulfate reaction /mol L-1s-1
19
1.67 x 10-5
1.67 x 10-5
8.00 6.00
26 39
1.22 x 10-5 8.13 x 10-6
1.22 x 10-5 8.13 x 10-6
10.00
5.00
48
6.60 x 10-6
6.60 x 10-6
5 6
10.00 10.00
3.00 10.00
86 21
3.69 x 10-6 1.51 x 10-5
3.69 x 10-6 1.51 x 10-5
7 8
8.00 6.00
10.00 10.00
27 38
1.17 x 10-5 8.34 x 10-6
1.17 x 10-5 8.34 x 10-6
9
5.00
10.00
45
7.04 x 10-6
7.04 x 10-6
10
3.00
10.00
82
3.87 x 10-6
3.87 x 10-6
Solution
Vol.S2O82/ mL
Vol. I- / mL
1
10.00
10.00
2 3
10.00 10.00
4
3)
Time/s
Calculation of the values for the logs of [I -], [S2O82-] and the rate R
Solution
[S2O82-]/M
[I-]/M
R /mol L-1s-1
log([S 2O82-]/M)
log([I -]/M)
log(R/Ms-1)
1
0.0385
0.0769
1.67 x 10-5
-1.42
-1.11
-4.77
2
0.0385
0.0615
1.22 x 10-5
-1.42
-1.21
-4.91
3
0.0385
0.0462
8.13 x 10-6
-1.42
-1.34
-5.09
4
0.0385
0.0385
6.60 x 10-6
-1.42
-1.42
-5.18
5
0.0385
0.0231
3.69 x 10-6
-1.42
-1.64
-5.43
6
0.0385
0.0769
1.51 x 10-5
-1.42
-1.11
-4.82
7
0.0308
0.0769
1.17 x 10-5
-1.51
-1.11
-4.93
8
0.0231
0.0769
8.34 x 10-6
-1.64
-1.11
-5.08
9
0.0192
0.0769
7.04 x 10-6
-1.72
-1.11
-5.15
10
0.0115
0.0769
3.87 x 10-6
-1.94
-1.11
-5.41
4)
Graph of log(R/Ms -1) against log([I -]/M)
Figure 1 Solution
log([I -]/M)
log(R/Ms -1)
1
-1.11
-4.77
2
-1.21
-4.91
3
-1.34
-5.09
4
-1.42
-5.18
5
-1.64
-5.43
Slope of best-fit line = 1.243
5)
Graph of log(R/Ms -1) against log([S 2O82-]/M) Solution
log([S 2O82-]/M)
log(R/Ms -1)
6
-1.42
-4.82
7 -1.51 Slope of best-fit line = 1.122 8 -1.64
-4.93 -5.08
9
-1.72
-5.15
10
-1.94
-5.41
Figure 2
6)
Experimental determination of the values, m and n in the rate equation and the calculation of the rate constant, k, for each solution
The rate law governing the reaction 2 I- (aq) + S2O82- (aq) → I2 (l) + 2 SO42- (aq) is given by R = k [I-] n [S2O82-] m Since [S2O82-] is held constant in solutions 1 to 5, R = k 1 [I-] n , where k 1 = k [S2O82-]m Hence, log R = log k 1 + n log [I-] and a plot of log R against log [I-] in Figure 1 gives a straight line of slope n. From Figure 1, Slope of best-fit line = 1.243 n = 1 (to the nearest integer)
Likewise, since [I-] is held constant in solutions 6 to 10, R = k 2 [S2O82-]m , where k 2 = k [I-] n Hence, log R = log k 2 + m log [S2O82-] And a plot of log R against log [S2O82-] in Figure 2 gives a straight line of slope m. From Figure 2, Slope of best-fit line = 1.122 m = 1 (to the nearest integer)
Calculation of k values for each solution
Since m = 1 and n = 1, k = R / ([S2O82-] [I-] )
Solution 1
[S2O82-]/M 0.0385
[I-]/M 0.0769
R /mol L-1s-1
2
0.0385
0.0615
5.15 x 10-3
3
0.0385
0.0462
1.22 x 10-5 8.13 x 10-6
4 5
0.0385 0.0385
0.0385 0.0231
6.60 x 10-6 3.69 x 10-6
4.45 x 10-3 4.15 x 10-3
6 7
0.0385 0.0308
0.0769 0.0769
1.51 x 10-5 1.17 x 10-5
5.10 x 10-3 4.94 x 10-3
8
0.0231
0.0769
8.34 x 10-6
4.69 x 10-3
9 10
0.0192 0.0115
0.0769 0.0769
7.04 x 10-6 3.87 x 10-6
4.77 x 10-3 4.38 x 10-3
7)
1.67 x 10-5
k /M-1s-1 5.64 x 10-3
4.57 x 10-3
Calculation of the mean of k and the mean deviation
k mean = (∑ Values of k) / 10 = 4.784 x 10-3 M-1s-1 = 4.78 x 10-3 M-1s-1 ( 3 sig.f) Deviation = k- k mean k mean / M-1s-1
Deviation (= k- k mean) / M-1s-1
Solution
k /M-1s-1
1
5.64 x 10-3
8.56 x 10-4
2
5.15 x 10-3
3.66 x 10-4
3
4.57 x 10-3
-2.14 x 10-4
4
4.45 x 10-3
-3.34 x 10-4
5
4.15 x 10-3
6
5.10 x 10-3
7
4.94 x 10-3
1.56 x 10-4
8
4.69 x 10-3
-0.94 x 10-4
9
4.77 x 10-3
-0.14 x 10-4
10
4.38 x 10-3
-4.04 x 10-4
-3
4.784 x 10
-6.34 x 10-4 3.16 x 10-4
∴ Mean deviation in k = (∑ Values of deviation) / 10
=0 8)
Graph of loge t against 1/T for the temperature-dependence experiment
Temperature
Temperature,T / K
Time,t /s
(1/T) / K -1
ln(t/s)
10.0
283.15
1144
3.53 x 10-3
7.04
20.0
293.15
322
3.41 x 10-3
5.77
30.0
303.15
86
3.30 x 10-3
4.45
45.0
318.15
16
3.14 x 10-3
2.77
60.0
333.15
4
3.00 x 10-3
1.39
o
/ C
Figure 3
Combining the equations derived from Arrhenius equation, we get loge t = Ea / RT + D Hence, the plot of loge t against 1/T in Figure 3 will give a straight line with slope Ea/R. From Figure 3, slope of best-fit line = 10735 ∴ Ea/R = 10735
Ea = 10735 x 8.314 J mol-1 K -1 = 89.3 kJ mol-1
Discussion
1)
Method of Initial Rates
When we use the method of initial reaction rates to determine the rate law of a
chemical reaction, we measure the initial rate of reaction at short instantaneous intervals before any significant changes in the concentrations of reactants occur. For us to determine the order of reaction with respect to each participating reactants, the initial concentration of that reactant is varied and the subsequent change in the rate of reaction observed. Such a method is useful as it saves time since complete reaction is not necessary in observing the initial rates. It is also a relatively efficient way of determining the rate law of a chemical reaction, especially for the study of slow reactions. However, its disadvantages include: 1. Requiring the need to perform multiple experiments with only the concentration of a single reactant varied, just to determine the order of reaction with respect to a that reactant. 2. Only working for slow reactions as, if the reaction proceeds too fast, the uncertainty as
well as percentage error in measurements would be significantly
high.
2)
Keeping the temperature for each reacting solution constant
Arrhenius proposed that for every reaction, the colliding molecules must exceed a minimum energy threshold called the activation energy, Ea, in order to react. At a given temperature, molecules possess a range of kinetic energies as shown by the Maxwell-Boltzmann distribution curve. Effective collisions would only occur between reacting molecules with energies greater than the activation energy required of the reaction. As the temperature is increased, the distribution curve shifts right (as shown in the figure below) resulting in a higher fraction of molecules having energies greater than Ea. Therefore, there is an increase in the number of effective collisions, leading to an increase in the rate of reaction.
In the first part of our experiment, where we seek to determine the order of reaction with respect to the reactants, varying concentrations of reactants shou ld be the only factor affecting the rate of reaction. By keeping the temperature constant for all the runs, it was ensured that the rate constant remained the same and any changes in the rate of reaction were solely due to changes in the initial concentrations of the reactants and hence, the order of reaction with respect to each reactant can be determined accurately.
3)
Keeping the total volume of the mixture constant
For our experiment, varying amounts of de-ionized water were added to each reacting solution so as to keep the total volume constant. Thus, the concentration of each reactant is proportional to the volume of reactant used. And so, by maintaining the volumes of all other reactants as constant, it is easy for us to vary the initial concentrations of the reactant simply by changing the volume of that reactant used. Essentially, the total volume was kept constant by adding de-ionized water to maintain the ionic strength of the reacting solution. The ionic strength, on a concentration basis, of a solution is a function of the concentration of all ions present in a solution. Ic = ½ Σ cBzB2 (where cB is the molar concentration of ion B (mol dm-3), z B is the charge number of that ion, and the sum is taken over all ions in the solution.)
In the study of solution reaction kinetics, it is more desirable to change the concentration of an active species without changing the ionic strength of the solution as the latter affects the rate constant by the kinetic ionic-strength effect . Since the charges of the reacting iodide and persulfate ions have the same sign, the activated complex would be much more charged than its reactants. Hence, increasing the ionic strength of the solution would increase the stability of the activated complex causing a decrease in the Ea and thus, increasing the rate constant. Our aim at the beginning of the experiment was to determine the rate law of the reaction, with all factors except one, being kept constant all the time. By maintaining the ionic strength of the solution, it ensured that the rate constant was kept unchanged and that any changes in the rate of reaction were solely due to changes in concentrations of the iodide and persulfate ion.
4)
Pseudo-1 st order / Isolation method
From the rate equation of the earlier experiment, R = k [I-] [S2O82-] we gather that 2 reactant molecules are involved in the rate determining step. By keeping the concentration of any one reactant constant (i.e. [S2O82-]) as the reaction progresses, the reaction can then be said to be pseudo-1st order with respect to the other reactant (i.e. [I-]) with new rate constant k obs = k [S2O82-] and now the rate of reaction only depends on changes in that reactant’s concentration (i.e. [I-]). The [S2O82-] should be set to be much higher than [I-] to ensure that changes in the concentration of S2O82- would not be significant throughout the course of the reaction.
5)
Precautions taken Determination of Reaction Orders and Rate Constant
i) The conical flasks used to hold the reacting solution were shaken as dry as possible between runs so as to ensure that the concentrations of the reactants are kept as accurate as possible since it has a direct effect on the rate of reaction.
ii) For each run, the contents were swirled for approximately the same length of time and then left to stand. This is to ensure that all reacting solutions were mixed
homogeneously and avoid significant changes between the rate of reactions caused by difference in swirling.
iii) Before the addition of the variable reactant for the reaction to occur, the temperature of the solution was taken and necessary alterations made to ensure that for all the runs, the reaction was occurring at a constant temperature. Any discrepancies in temperature would affect the rate constant.
iv) When timing, the stopwatch was only started when the last drop of reactant in the test tube was added to the reactant in the conical flask. This minimizes lapses in timings.
v) The conical flasks of the solutions were placed on a white piece of paper in every run so that any colour changes would be distinct and observed instantaneously.
Determination of the Temperature Effect on a Chemical Reaction
i) The temperature of the water bath was constantly checked to ensure that it is kept almost constant. For the runs carried out between 0 oC to 20 oC, any change in the temperature of the water bath was promptly reversed by adding ice or water.
ii) During the course of the reaction, the reacting mixtures were fully immersed in the water bath to ensure that all parts of the mixture are at constant temperature throughout the experiment. Any slight changes in temperature would result in significant changes in the reaction rates and thus, affect the accuracy of results.
iii) To accurately determine when the reacting mixture turns colourless, a flask of colourless reactant was placed next to the reacting mixture for simple comparison.
iv) More than one set of timings were taken for each run so as to reduce the effect of human error.
Conclusion
Determination of Reaction Orders and Rate Constant Experimental results have shown that the order of reaction with respect to I- and S2O82for the reaction, 2 I- (aq) + S2O82- (aq) → I2 (l) + 2 SO42- (aq) are both 1 respectively. Thus, it is an overall second order reaction. Using these values, the average rate constant of the reaction, k mean, was calculated to be 4.78 x 10-3 M-1s-1. Since the mean deviation of the value k is 0, we conclude that the method of initial reaction rates is a relatively
accurate approach to determine the reaction orders, rate constants and hence, the rate law for the reaction between iodide and persulfate ions.
Determination of the Temperature Effect on a Chemical Reaction From our experimental results, the activation energy, Ea of the reaction between KMnO4 and oxalic acid was found to be 89.3 kJ mol-1. We can infer that temperature is inversely proportional to the time needed for the reaction to be completed. Hence, we conclude that temperature is linearly related to the rate of reaction and this agrees with the relation offered by the Arrhenius equation.
References
P.W. Atkins & J.de. Paula, 2006. Atkins’ Physical Chemistry 8th Ed. Published by Oxford, Portland. M.S. Silberberg, 2003. Chemistry the Molecular Nature of Matter and Change 3rd Edition. Published by McGraw Hill.
Chemical Kinetics: Method of Initial Rates Retrieved 18th February 2008 from the World Wide Web: http://www.chm.davidson.edu/ChemistryApplets/kinetics/MethodOfInitialRates.html Chemical Kinetics: Isolation Method Retrieved 18th February 2008 from the World Wide Web: http://www.chm.davidson.edu/ChemistryApplets/kinetics/IsolationMethod.html Wikipedia – Ionic Strength Retrieved 18th February 2008 from the World Wide Web: http://en.wikipedia.org/wiki/Ionic_strength Kinetic electrolyte effect Retrieved 18th February 2008 from the World Wide Web: http://www.iupac.org/goldbook/K03401.pdf
Kinetics of the Persulfate-iodide Clock Reaction Retrieved 18th February 2008 from the World Wide Web: http://ptcl.chem.ox.ac.uk/~hmc/tlab/experiments/502.pdf.
