Lab 2- Iodine Clock Reaction PDF

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For Dr. Lu...

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Lab 2: A Kinetic Study of an Iodine Clock Reaction By Hayden Casassa Partner: Jonathan Baugh 2/22/16

Learning Objectives: 1. To learn about the energy of activation for the reaction. 2. To determine K (the rate constant), of the same rate law. 3. To learn about the order of the rate law (m and n) for the iodine clock reaction using initial rates.

Data and Results: Table 1: Experiment 1 performed at 296.3 K Experiment

Initial [I-]

Initial [S2O82-]

Elapsed Time (s)

Reaction Rate (M/s)

1a

0.0400 M

0.0400M

62 s

16 M/s

1b

0.0400 M

0.0400M

51 s

20 M/s

2a

0.0800 M

0.0400M

45 s

22 M/s

2b

0.0800 M

0.0400M

46 s

22 M/s

3a

0.0800 M

0.0200 M

93 s

11 M/s

3b

0.0800 M

0.0200 M

95 s

11 M/s

Average Rate (M/s)

18 M/s

22 M/s

11 M/s

Order of reaction with respect to I- ____0.29__________ Order of reaction with respect to S2O82- _____1.0_______ Overall order of Reaction _______1.29_________ Table 2: Experiment 2 performed at 4 different temperatures Experiment

Temp (K)

1/T

Elapsed Time (s)

2 296.3 K

0.0034 45.4 s

4 285.5 K

k

ln k 1,100

7.0

0.0035 97 s

520

6.3

5 304.3 K

0.0033 67 s

780

6.7

6 317.5 K

0.0031 32 s

1600

7.4

Slope of Trend line: -2342.9 Calculated value of Ea (KJ/mol): 19.47 kJ/mol

Sample Calculations: Initial [I-]: Mstock*Vstock=Mmixture*Vmixture (0.200M)*(1.00 ml)=(x)*(5.00 ml) x=0.0400 M Initial [S2O82-]: Mstock*Vstock=Mmixture*Vmixture (0.100M)*(2.00 ml)=(x)*(5.00 ml) x=0.0400 Order with Respect to I22 M/s = (0.0800M)) ^ x 18 M/s (0.0400M)) x=0.29 Order with Respect to S2O8222 M/s = (0.0400M)) ^x 11 M/s (0.0200M)) x=1 Solving for K 22 M/s =k[0.0800M]0.29 * [0.0400M] k= 1,100

Solving for Activation Energy M=-Ea/8.314 -2342.9=Ea/8.314 Ea= 19.47 kJ/mol Converting from Celsius to Kelvin oK=273 +C oK=273 +31.3oC oK=304.3 Discussion: The Iodine Clock Reaction lab is a lab that focus on the rate in which a reaction takes place. the Rate Law equation can be determined by taking the concentration of the reactants and multiplying them by their order of reaction. The order of reaction can be determined from dividing the rate of reaction y by the rate of another (x) and setting it equal to the concentration of reaction y divided by the concentration of reaction x. Then by raising the side with the concentrations to an unknown number (typically m or n), you can solve for them using logarithms in order to determine the order of the reaction for the given concentration. Finally, buy multiplying the concentrations by constant K it produces the rate law equation which is Rate=k[A]m[B]n. The reaction studied in the lab between the persulfate ion and the iodide ion was 2I- + S2O8 —> I2 + 2SO4 . This reaction allowed us to collect convenient measurements for the rate equation but instead, known amounts of S2O3 were added to the mixture because it was determined that it would react instantly with the Iodine (I2). When the reaction finished, any additional Iodine present in the solution would react with the starch indicator to form a blue color signifying that the second reaction reached its equilibrium point. After determining the concentrations of the reactants and observing the reaction rates, it was determined that the Order of Reaction with respect to I- and S2O82- is 0.29 and 1.0 respectively. Therefore, the overall order of reaction is 1.29. We determined that for experiments 1 and 2 we could calculate the order of reaction for I- since the persulfate concentration remained the same at 0.0400 M while I- changed from 0.0400 M to 0.0800 M. Experiments 2 and 3 were used to determine the order of reaction for persulfate as the concentration of iodine stayed constant at 0.0800 while the concentration of the persulfate ion changed from 0.0400 M to 0.0200 M. The precision in experiment one was very accurate in trials 2 and 3 as they were only 1-2 seconds off from each other. However, experiment 1 was 10 seconds off. In Experiment 2, the data was skewed from the start of the experiment. The trend that was hypothesized that at a constant concentration as the temperature increased the time of the reaction would decrease. However, at 296.3 K the elapsed time was 45.5 seconds while at 306.7 K the reaction took 68 seconds. Both trials were performed twice but the results were only one second apart from the initial trial in each one. This impacted our trend line and our K results which were found using the rate law from the concentrations in part 1. The trend line did decrease which supported the hypothesis but the accuracy in the slope of the line is questionable in its accuracy. This would

subsequently impact the activation energy which was determined from the line to be 19.47 kJ/ mol by setting the slope equal to -Ea/8.314. One source of error in the experiment could be the inhomogeneity of the solution. If the reactants weren't mixed well enough with the starch indicator then the time of the reaction could be off from the actual value. Also, another possible source of error could be the uncertainty of the concentration of the chemicals given due to the inconsistency of the time in comparison to temperature. However, in order to avoid human sources of error, any experiment with questionable results were duplicated in order to confirm that the results were accurate and precise....