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Kinetics Lab Report
Comprehensive General Chemistry 3 (CHEM 11300)
University of Chicago
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Kinetics: "Iodine Clock" Lab Report
Introduction: In order to describe the chemical kinetics of a reaction, it is desirable to determine how the rate of reaction varies as the reaction progresses. The rate law is a mathematical equation that describes the progress of the reaction and has the following general form for the reaction aA + bB --> cC + dD:
rate = -d[A]/(a(dt)) = k[A]m[B]n
[A] and [B] are concentrations of the reactants and k is the rate constant. The variables m , n and k can only be determined experimentally. In this experiment, the oxidation reaction of iodide by peroxydisulfate is performed. The rate law for this reaction is as follows:
rate = -Δ[S 2 O 8 2-]/Δt = k[I-]m[S 2 O 8 2-]n
This lab provides an opportunity to understand different concepts of chemical kinetics such as the reaction rate, rate constant, and reaction order. In this lab-- using several mixtures of the iodide and peroxydisulfate solutions-- it is possible to calculate the reaction order and the reaction constant of the chemical reaction. By manipulation of temperature, it is also possible to understand the temperature dependence of the rate constant as well as determine the activation energy of the reaction using an Arrhenius plot. Understanding chemical kinetics is important because of its practical applications. For example, understanding rates of reaction is important in the manufacturing of ammonia for fertilizer using the Haber process. If the reaction to produce ammonia is conducted at a low temperature, the rate of reaction will be too slow. Because the manufacturers want to produce as much fertilizer as possible, they must manipulate the temperature to produce a satisfactory rate.
Experimental: For the first portion of this experiment, 5 "Mixtures" are prepared, with varying concentrations of solutions; the concentrations are summarized in the following table:
Mixture # I- (0) SO 4 2- (0)
S 2 O 3 2- (0) Starch (drops) H 2 O S 2 O 8 2-(0)*
1 50 25 10 10 5 10
2 25 30 10 10 25 10
3 10 33 10 10 37 10
4 10 27 10 10 28 25
5 10 16 10 10 14 50
All of the solutions amounts are in mL, except starch. Combine all of these solutions, except S 2 O 8 2- into a beaker and bring the Mixture to room temperature in a water bath. Separately heat S 2 O 8 2- in the same bath. Once the Mixture and S 2 O 8 2- reach the same temperature, combine the two while on a magnetic stir plate and begin recording the time on a stopwatch. After stirring for ten seconds, return the combined Mixture to the bath and stop the stopwatch once the Mixture reaches a blue color.
For the second portion of this experiment, repeat the above procedure with three volumes of Mixture 4. Run the same experiment, but in a 10 °C, 30 °C, and 40 °C bath for trials 1, 2, and 3, respectively (we don't need to do 20 °C because our room temperature trial in the first part of the experiment serves as our data point).
Record the initial and final temperature for every trial; take note of qualitative observations at each step and dispose of each solution in the Iodine waste bin.
[No deviations were made from UChicago General Chemistry Lab Manual]
Data Analysis:
Qualitative
For all mixtures, once the peroxydisulfate was added the solutions turned from clear to blue after a certain amount of time Δt.
Quantitative
Table 1: Rate of Reaction for Each Mixture
Mixture # Tinitial (°C) Tfinal (°C) Δt (s) -Δ[S 2 O 8 2-]/Δt (M/s)
1 19 19 83 6-
2 19 19 158 3-
3 19 19 398 1-
4 19 19 202 2-
5 19 19 97 5-
Sample Calculation; -Δ[S 2 O 8 2-]/Δt, Mixture 1:
Moles of S 2 O 3 2- = 0 x .01M = .0001 mol S 2 O 3 2- x (1 mol S 2 O 8 2-/1 mol S 2 O 3 2-) = .00005 mol S 2 O 8 2-/0 = .0005M S 2 O 8 2- .0005M S 2 O 8 2-/83s = 6-6 M/s
0 x 0 = 0 mol/0 = .035 mol S 2 O 8 2- log(.035) = -1.
Figure 2: Relationship Between Peroxydisulfate Concentration and Rate
Table 4: Raw Data for Each Temperature
Target Temperature (°C)
Tinitial (°C) Tfinal (°C) Δt (s)
10 11 11 253
20 19 19 202
30 27 27 105
40 39 39 43
Table 5: Relationship Between Temperature and Rate
Target T (°C) Average T (°C) Δt (s) -Δ[S 2 O 8 2-]/Δt (M/s) k(T)
10 11 253 1-3 1-
20 19 202 2-6 2-
30 27 105 4-6 3-
40 39 43 8-6 7-
Sample Calculation; m and kb: log(-Δ[S 2 O 8 2-]/Δt) = log(kb) + m x log([I-]) From Figure 1... y = 0 - 4. m = 0. kb = 10-4 = 5-
Sample Calculation; n and ka: log(-Δ[S 2 O 8 2-]/Δt) = log(ka) + n x log([S 2 O 8 2-]) From Figure 2... y = 0 - 4. n = 0. ka = 10-4 = 4-
Sample Calculation; k: ka = k[I-]m k = (4-5)/(00) k = 2-
kb = k[S 2 O 8 2-]n k = (5-5)/(00) k = 2-
Average k = 2-
Sample Calculation; k(T), 10 °C: rate = -Δ[S 2 O 8 2-]/Δt = k[I-]m[S 2 O 8 2-]n 1-6 = k(0)0(0)0. k(10 °C) = 1-
Table 6: Different Logarithms of Various Rate Constants
T (°C) T (K) 1/T (K-1) k(T) log(k) ln(k)
11 284 3-3 1-3 -2 -6.
temperatures for Mixture 4, it was possible to calculate the activation energy. Using the Arrhenius plot (figure 3) and the equation, ln(K) = ln(A) - (Ea/R)(1/T), the activation energy was computed to be 41 kJ/mol.
There are several possible sources of error that could affect our rate constant and activation energy calculations. Determining when the reaction was complete was problematic. It was uncertain whether the first presence of blue signified the end of the reaction or the presence of blue throughout the solution signified the end. This discrepancy affects the change in time, which affects the calculation both the rate constant and the activation energy. We waited until the entire solution was blue, which could have consistently overestimated our reaction times and consequently underestimated our rates. Another possible source was the variation of temperature for each sample. In order to produce accurate rate constants, the temperature must remain constant for the entire reaction. It is possible to keep the temperature constant by observing the temperature and adding ice or boiling water to the bath as needed. The presence of unwanted ions in solutions can also have an affect on calculations. The presence of metal ions such as copper, even at small concentrations, can catalyze the reaction. In each mixture, a different concentration of tap water was added with different concentrations of unwanted ions. This error can be avoided by adding a complexing agent like sulfate and using distilled water, which is available but unreliable in our lab.
Considering the consequences of the kinetic salt effect, sulfate solution was added to all mixtures in the experiment. The kinetic salt effect is effect of the salts in solution on the rate of reaction. When reactions involving ionic compounds occur in solution, the addition of a salt can speed up, slow down or have no effect on the rate of reaction. Depending on the preexisting ionic strength of the solution, a certain amount of sulfate solution is added to each mixture to keep constant the ionic strength for all mixtures. In Mixture 1, the moles of sulfate were 0 x .33M = 0 moles. In mixture 2, 3, 4, and 5 the moles of sulfate were 0, 0, 0, and .00528, respectively.
The bath temperature changes only slightly because of the high specific heat of water. The equation q = mcΔT illustrates the reasons why it is difficult to effect the temperature significantly. Because the mass of the bath is large and the specific heat of water is relatively large, the samples would need a large amount of heat energy to produce a large change in temperature.
Conclusion: This lab provided an opportunity to experimentally determine the rate constant and the order of the reaction in which iodide was oxidized by peroxydisulfate reaction. In this lab, the rate constant was determined to be .00227 L mol-1 s-1 and the total order of the reaction was
- The overall rate expression was rate = [I-]0[S 2 O 8 2-]0. The "correct' values for the m and n of this reaction are approximately 1; we got quite close to these values. The effect of temperature on the rate of constant was also examined in this lab. Plotting these effects produced an Arrhenius plot that was used to derive the activation energy for the reaction of 41. 6 kJ/mol. There were some sources of error like the endpoint of the reaction, the variations in temperature, and the presence of interfering ions. It was difficult to determine which shade of
blue signified the end of the reaction. It was also difficult to keep the temperature constant for the temperature portion of the lab. Having a supply of ice and boiling water could aid in minimizing variations. The variations of the temperature were not an issue for the room temperature portion because the bath water had a large mass and the specific heat of water is relatively large. Finally, the presence of ions like copper could catalyze the reaction. This unwanted effect could be reduced by using a complexing agent or by using distilled water. Another source of error was the effect of ionic strength on the rate of reaction. The sulfate was added to keep the ionic strength equal for all mixtures.
Sources Cited: 1. Zhao, Meishan and Dragisich, Vera. General Chemistry Experiments. Hayden-McNeil Macmillan Learning, 2018.
Kinetics Lab Report
Course: Comprehensive General Chemistry 3 (CHEM 11300)
University: University of Chicago
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