In the Laboratory
Kinetics of Platinum-Catalyzed Decomposition of Hydrogen Peroxide
W
Tiffany A. Vetter and D. Philip Colombo, Jr.* Department of Chemistry, Rockhurst University, Kansas City, MO 64110; *
[email protected] Several methods to increase student interest and learning in general chemistry laboratory have been proposed (1). Many of these involve student control over the procedure used to investigate chemical systems and experiments with realworld applications (2–5). It is difficult for first-year chemistry students to independently generate adequate experimental procedures to investigate some general chemistry concepts (6). Kinetics, for example, may require established general chemistry techniques to explore. We propose a simple experiment where students employ a detailed procedure to generate kinetic data using a consumer product to determine safety information pertinent to product use. Overview CIBA Vision Corporation markets a contact lens cleaning system1 that consists of an AOSEPT disinfectant solution and an AOSEPT lens cup. The disinfectant is a buffered 3.0% m兾v hydrogen peroxide solution and the cup includes a platinum-coated AOSEPT disc. The hydrogen peroxide disinfects by killing bacteria, fungi, and viruses found on the contact lenses (7, 8). Because the concentration of hydrogen peroxide needed to disinfect is irritating to the eyes, the hydrogen peroxide must be neutralized or decomposed before the contact lenses can be used again. Following the kinetics of the decomposition of hydrogen peroxide by measuring the amount of oxygen generated is a common general chemistry experience (9–17). This experiment involves investigating the heterogeneously catalyzed decomposition of hydrogen peroxide: 2H2O2(l)
Pt
2H2O(l) + O2(g)
The AOSEPT disc used as a catalyst in this experiment is platinum black coated on a plastic support and is easily obtained from any drugstore for about five dollars. The observation of the reaction kinetics requires minimal solution preparation and is easily followed. The reaction kinetics are simplified in that only one of the reactant concentrations can vary. Finally, students gain experience using integrated rate laws by calculating the time one must wait until the contact lenses can be safely used after disinfection. Experiment Solutions of varying hydrogen peroxide concentrations are prepared from a standard 3% solution. These solutions are buffered with phosphates to an approximate pH of 7. An AOSEPT disc is prepared by soaking it for 30 minutes in a hydrogen peroxide solution that is similar to the one being studied. Once the disc is initially prepared, subsequent soakings can be shortened to a few minutes. After the solution and the disc are prepared, the disc is placed in solution 788
and the amount of generated oxygen is measured as a function of time. The initial rate of the reaction is determined from the plot of the amount of oxygen versus time. Either the change in volume of generated oxygen under conditions of constant pressure can be measured with a leveling bulb (15, 16) or the change in pressure of generated oxygen under conditions of constant volume can be measured with a computerized, low-pressure sensor (17). Both methods yield similar results. The order of the reaction with respect to the hydrogen peroxide and the rate constant are determined by comparing the rate of the reaction for a series of different hydrogen peroxide concentrations. The energy of activation of the reaction is determined by comparing the rate of the reaction at a series of different temperatures. The integrated rate law is used to determine the time required to decompose the hydrogen peroxide to a concentration that is safe for eyes. Hazards Contact with hydrogen peroxide can irritate the skin and eyes. Wear safety goggles and avoid contact with skin. If skin or eye exposure occurs, flush with copious amounts of water. Calculations The plots of generated oxygen versus time are linear over the first few minutes. These initial rates are converted from units of milliliters per second or kilopascals per second to moles of oxygen generated per second by using the ideal gas law. The rate of oxygen formation is converted to the rate of hydrogen peroxide decomposition, in units of moles of hydrogen peroxide per second, using the reaction stoichiometry. This rate is converted to a rate with units of moles per liter of hydrogen peroxide per second by dividing the amount of hydrogen peroxide by the volume of solution. The students determine the order of the reaction from the slope of the plot of the natural log of the rate versus the natural log of hydrogen peroxide concentration (Figure 1). One could also confirm that the rate law is first order with respect to the hydrogen peroxide from the linear plot of rate versus hydrogen peroxide concentration. Once the order of the reaction is determined, the rate constant is found to be 6 × 10᎑4 ± 2 × 10᎑4 s᎑1. The rate constant at a series of temperatures is used to determine that the energy of activation of the catalyzed reaction is 40 ± 10 kJ兾mol (Figure 2). Results Both the rate constant and the energy of activation vary from disc to disc and for a given disc from day to day. The rate constant and the integrated form of the rate law are used to determine the time required to reduce the concentration
Journal of Chemical Education • Vol. 80 No. 7 July 2003 • JChemEd.chem.wisc.edu
In the Laboratory -7.0
-8.0
y = -5004.4 x + 9.1493 R 2 = .9966 -8.5
ln(k)
ln(Rate)
-7.5
-9.0
-8.0
y = 1.00x − 8.08
-9.5
R 2 = .99 -8.5 -10.0 -1.5
-1.0
-0.5
0.0
3.2
3.3
T
ln[H2O2]
3.4
ⴚ1
/ (10
ⴚ3
3.5
K
ⴚ1
)
Figure 1. Plot of the natural log of the rate constant versus the natural log of the hydrogen peroxide concentration. The slope of this line is the order of the reaction with respect to the hydrogen peroxide. These data are from the low-pressure sensor method.
Figure 2. An Arrhenius plot of the natural log of the rate constant versus the reciprocal of the temperature in Kelvin. The energy of activation, which is calculated by the linear regression of these data, is 42 kJ/mol.
of hydrogen peroxide from 3.0% m兾v to 60 ppm, the irritation threshold. The experimentally determined time is approximately 3 ± 1 hours: this is half of the 6 hours CIBA recommends.
ide solution for at least six hours. Below the baskets, but attached to the lid, is the platinum-coated disc. Not only does the disc catalyze the decomposition of the hydrogen peroxide, but its placement below the basket containing the contact lenses allows the bubbles of generated oxygen to aid in the removal of surface proteins. The entire cup and lid are replaced when additional solution is purchased.
Conclusion Student response to the experiment has been positive. Evaluations showed that students enjoyed the real-world nature of the experiment and were motivated to use kinetics to solve the problem. Although variation in the rate constant and the energy of activation from group to group exists, almost all students successfully determine that the reaction is first order with respect to the hydrogen peroxide and calculate a safe reaction time that is less than the six hours recommended by CIBA. WSupplemental
Material
Instructions for the students, including background kinetic information, and detailed notes for the instructor, including examples of experimental data and graphs and postlab questions, are available in this issue of JCE Online. Acknowledgments The authors wish to thank Barbara Heyl of CIBA Vision Corporation for the generous donation of AOSEPT discs and helpful conversations and Father James Wheeler for attempting this experiment in his general chemistry class. This work was made possible with the generous support of Rockhurst University through funding of a Summer Research Fellowship for Tiffany Vetter and of a Presidential Research Grant for Philip Colombo. Note 1. The contact lenses are rinsed, placed in small baskets attached to the lid of the cup, and lowered into the hydrogen perox-
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JChemEd.chem.wisc.edu • Vol. 80 No. 7 July 2003 • Journal of Chemical Education
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