In the Laboratory
Wash Bottle Laboratory Exercises: Iodide-Catalyzed H2O2 Decomposition Reaction Kinetics Using the Initial Rate Approach Rebecca Barlag* and Frazier Nyasulu Department of Chemistry and Biochemistry, Ohio University, Athens, Ohio 45701 *
[email protected] The catalyzed H2O2 decomposition reaction is catalyst
2H2 O2 ðaqÞ sf 2H2 OðlÞ þ O2 ðgÞ
(1Þ
A procedure to determine the kinetics of the iodide-catalyzed H2O2 decomposition reaction based on water displacement from a wash bottle has been reported by Teggins and Mahaffy (1). In this procedure, the H2O2 and KI reagents totaling 300 mL are mixed and transferred to a wash bottle. As O2(g) is generated, some of the reaction solution is displaced. The initial rate of reaction is determined by measuring the time to displace specific volumes. By comparing the initial rates for different runs, the orders of reaction were both found to be 1. Compared to conventional water displacement schemes (2, 3), this approach is attractive because it is easily performed. We have reported a wash bottle water displacement scheme to determine the % NaHCO3 in an Alka-Seltzer tablet, the molar mass of CO2, and the ideal gas constant R (4). Encouraged by the quality of results obtained by students, we have examined the Teggins and Mahaffy (1) wash bottle water displacement scheme and found the following changes to be helpful: (i) Reagents are added to a test tube placed in a water-filled wash bottle so that water is displaced instead of the reaction H2O2-KI solution. (ii) Small volumes of reagents totaling 5 mL are used. Theory
It is essential to include V(solution) in the definition of initial rate in eq 2 so that the rate constant in eq 4 is applicable to any volume of reaction solution. Initial Rate Measurement Based on Concentration of H2O2 Based on H2O2 concentration, the initial rate is Δ½H2 O2 initial rate ¼ 2Δt
Lab measurements will provide V(O2(g))/(V(solution) Δt) which will be converted to -Δ[H2O2]/(2Δt) as follows: V ðO2 ðgÞÞ mol O2 f V ðsolutionÞ Δt V ðsolutionÞ Δt f
mol H2 O2 reacted Δ½H2 O2 f V ðsolutionÞ Δt 2Δt
(6Þ
The temperature of the gas in the wash bottle and the room atmospheric pressure (P(atm) = P(O2(g)) are measured so that the V(O2(g)) can be converted to mol O2(g) according to the ideal gas law. Considering the equation Δ½H2 O2 (7Þ initial rate ¼ k½H2 O2 h ½I - i 2Δt the rate constant will have units of rate constant units ¼ Mð1 - h - iÞ s - 1
Initial Rate Measurement Based on Volume of O2(g) The initial rate is defined as V ðO2 ðgÞÞ initial rate ¼ V ðsolutionÞ Δt
(2Þ
where V(O2(g)) is the volume of the O2(g) in liters (L), V(solution) is the volume of the reaction solution in liters (L), and Δt is the displaced water collection time in seconds (s). Considering the equation V ðO2 ðgÞÞ (3Þ ¼ k½H2 O2 h ½I - i initial rate V ðsolutionÞ Δt where k is the rate constant, h is the order with respect to H2O2, and i is the order with respect to I-, the rate constant will have units of LðO2 ðgÞÞ rate contant units ¼ LðsolutionÞ M hþi s
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(4Þ
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(8Þ
The rate constant for the H2O2 decomposition is conventionally expressed according to eq 8. Order of Reaction with Respect to H2O2 With [I-] fixed and the [H2O2] varied from one run to another Δ½H2 O2 ¼ k½I - i ½H2 O2 h initial rate ¼ 2Δt ¼ k 0 ½H2 O2 h
(9Þ
where k 0 ¼ k½I - i
(10Þ
The order of the reaction with respect to [H2O2] can be
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In the Laboratory Table 1. Run Composition for Determining the Order with Respect to H2O2 Run No.
V(15.0% H2O2)/mL
V(0.500 M KI)/mL
V(H2O)/mL
1
0.00
1.00
4.00
2
0.50
1.00
3.50
3
1.00
1.00
3.00
4
1.25
1.00
2.75
5
1.50
1.00
2.50
Table 2. Run Composition for Determining the Order with Respect to KI V(15.0% H2O2)/mL
V(0.500 M KI)/mL
V(H2O)/mL
6
1.00
0.00
4.00
7
1.00
1.00
3.00
8
1.00
2.00
2.00
9
1.00
3.00
1.00
10
1.00
4.00
0.00
Run No. Figure 1. Wash bottle water displacement scheme.
determined by taking natural logarithms of eq 9 leading to the equation (11Þ lnðinitial rateÞ ¼ h ln½H2 O2 þ lnðk 0 Þ A plot of ln(initial rate) versus ln[H2O2] has a slope equal to of h, which is the order with respect to [H2O2], and a y-intercept ln(k0 ). Rather than implementing the ln versus ln plot, it is more instructive to determine the order of reaction via hypothesis testing. The hypotheses are the guessed orders of reaction with respect to [H2O2], namely h = 1, 2, 3, and so forth. The hypothesis that h = 1 is tested by plotting the initial rate versus [H2O2]. This is valid if the plot is linear and goes through the origin; the slope is equal to k0 . To test the hypothesis that h = 2, initial rate versus [H2O2]2 is plotted. Order of Reaction with Respect to IWhen the reaction is run with [H2O2] fixed and [I-] varied Δ½H2 O2 ¼ k½H2 O2 h ½I - i ¼ k 00 ½I - i (12Þ initial rate ¼ 2Δt where k 00 ¼ k½H2 O2 h
Run No. V(O2(g))/L Time/s [H2O2]/M
Rate (-Δ[H2O2]/ (2Δt))/M-1 s-1
1
0.00000
NA
0.000
0.000
2
0.00400
66
0.469
0.000
3
0.00831
64
0.939
0.001
4
0.01093
68
1.174
0.001
5
0.01522
62
1.878
0.002
first drops into the test tube, and the mass of the water collected over ∼60 s time period is determined. The temperature of the air in the wash bottle is measured by opening the lid just enough to insert a thermometer. The room pressure, the density of the water, and the density of the 15.0% H2O2 are also determined.
(13Þ
The order of the reaction with respect to [I-] is determined by the same hypothesis testing. Rate Constant (k) Having determined the orders of reaction, the rate constant is calculated on the basis of eqs 3, 9, 10, 12, and 13. Experimental Section Materials and Equipment 500-mL wash bottle, thermometer, barometer, 6-in. test tubes, milligram mass balance, digital pipettor, stir bar, stir plate, timer, large beaker, 15.0% H2O2, and 0.500 M KI. Procedure A photo of the setup is shown in Figure 1. A test tube containing a stir bar is placed in a wash bottle that is half-filled with water, and the run composition indicated in Table 1 (or Table 2) is added. The wash bottle lid is then closed, and a preweighed test tube is held to catch the water displaced from the wash bottle. The timer is started when water
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Table 3. Effect of [H2O2] on the Initial Rate of Reaction
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Hazards The H2O2 solution is an oxidizer and should be handled with care. Results and Discussion Reaction Order with Respect to H2O2 The initial rates as reported by a student participant are shown in Table 3 where the concentration of KI is fixed at 0.100 M and the concentration of H2O2 is varied from 0.00 to 1.88 M. To test the hypothesis that the order of reaction with respect to H2O2 is 1 (h = 1), a plot of initial rate (Δ[H2O2]/(2Δt)) vs H2O2 concentration is plotted; see Figure 2. Initial rate is directly proportional to concentration with the equation (14Þ y ¼ 0:00108x - 0:00001 R2 ¼ 0:9982 Equation 14 verifies that the order of the reaction is first order with respect to H2O2. Reaction Order with Respect to IThe initial rates as reported by a student participant are shown in Table 4 where the concentration of H2O2 is fixed at
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In the Laboratory Table 5. Rate Constants Equation Used To Calculate k
0.0109 ( 0.0004
eq 9
Figure 2. Effect of the H2O2 concentration on the initial reaction rate.
Table 4. Effect of [KI] on the Initial Rate of Reaction V(O2(g))/L
Time/s
[KI]/M
Rate (-Δ[H2O2]/(2Δt))/ M-1 s-1
6
0.0000
NA
0.000
0.00000
7
0.0079
61
0.100
0.00106
8
0.0153
60
0.200
0.00209
9
0.0231
60
0.300
0.00315
10
0.0219
40
0.400
0.00448
Run No.
k Value/M-1 s-1
eq 10 using results in eq 14
0.0108 ( 0.0003
eq 12
0.0114 ( 0.0004
eq 13 using results in eq 15
0.0118 ( 0.0003
reported by students based on eq 9 is 0.0112 ( 0.0002 M-1 s-1(N = 17) at ∼20 °C. The literature values are in the range 0.0100-0.0118 M-1 s-1 (5, 6). The average rate constant based on eq 3 (N = 13) is LðO2 ðgÞÞ 0:275 ( 0:002 LðsolutionÞ M 2 s
Conclusion This lab is instructive for the following reasons: (i) It is based on very basic equipment, and the procedure is straightforward and easy to implement and uses small volumes of reagents. (ii) It incorporates hypothesis testing. (iii) By calculating the rate constant in several ways, students gain a better understanding of the relationships between the various equations. (iv) It requires considerable data analysis in the Excel spreadsheet. (v) Results obtained compare well with literature values. We have written this lab at an advanced level. This lab can be simplified by excluding some parts. For example, the orders of reaction can be found by comparing reaction rates in Tables 3 and 4, and the rate constant can be calculated by using only one of the suggested equations. Literature Cited
Figure 3. Effect of the KI concentration on the initial reaction rate.
Equation 15 verifies that the reaction is first order with respect to [KI].
1. Teggins, J.; Mahaffy, C. J. Chem. Educ. 1997, 75, 566. 2. Nelson, J. H.; Kemp, K. C. Laboratory Experiments, Chemistry, The Central Science, 9th ed.; Prentice Hall: Upper Saddle River, NJ, 2003; pp 325-338. 3. Milio, F. R.; Debye, N. W. G.; Metz, C. Experiments in General Chemistry; Saunders College Publishing: Philadephia, 1991; pp 338-348. 4. Nyasulu, F.; Paris, S.; Barlag, R. J. Chem. Educ. 2009, 86, 842844. 5. Hansen, J. C. J. Chem. Educ. 1996, 73, 728–732. 6. Liebhafsky, H. A. J. Am. Chem. Soc. 1932, 54, 1792–1805.
Determination of the Rate Constant (k)
Supporting Information Available
0.939 M and the concentration of I- is varied from 0.000 to 0.400 M. To test the hypothesis that the order of reaction with respect to [I-] is 1 (i = 1), the initial rate (Δ[H2O2]/2Δt) vs [KI] is plotted (see Figure 3). Initial rate is directly proportional to concentration with the equation (15Þ y ¼ 0:01105x - 0:00005 R2 ¼ 0:9974
-1 -1
The calculated rate constants in units of M s obtained by a student participant are summarized in Table 5. The average
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Instructor notes, lab write-up, and postlab spreadsheet with sample data. This material is available via the Internet at http://pubs.acs.org.
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