In the Laboratory edited by
The Microscale Laboratory
Arden P. Zipp SUNY–Cortland Cortland, NY 13045
Small-Scale Kinetic Study of the Catalyzed Decomposition of Hydrogen Peroxide Ronald O. Ragsdale* and Jan C. Vanderhooft Department of Chemistry, University of Utah, Salt Lake City, UT 84112 Arden P. Zipp Chemistry Department, State University College at Cortland, Cortland, NY 13045 The rates of chemical reactions and the factors that affect them are important topics in most general chemistry courses. Many experiments show qualitatively the effect of concentration, temperature, and a catalyst on the rates of chemical reactions, but it is difficult to obtain quantitative data in a general chemistry laboratory. The direct chemical analysis of the concentration of a reactant or product versus time is desirable but frequently very labor intensive. Indirect physical methods are easily applied in advanced courses, but often only limited instrumentation is available in beginning courses. The iodine clock reaction involving the oxidation of iodide ion by peroxydisulfate ions or hydrogen peroxide has been a popular experiment, but it suffers from the use and manipulation of many chemicals. In addition, the complexity of the reaction and the data analysis required may confuse students. We report a reaction that can be studied directly and quickly by determining the rate of formation of oxygen bubbles produced according to the equation 2H2O2(aq) → 2H2O(l) + O2(g) The chemical costs are very low and there are no chemical disposal problems. The catalyst, a piece of pyrolusite rock, is recycled. The Experiment This experiment was first reported as a microscale lab (1) requiring only a 24-well plate, a graduated 2-mL pipet, hydrogen peroxide, a piece of pyrolusite, and a number 2 one-hole rubber stopper. There are several problems in doing this experiment on a microscale basis. First, a large percentage of the small amount of the H2 O2 decomposes in the time required to collect data, so that good initial rate data cannot be obtained. Second, the amount of heat released can affect the temperature of the more concentrated H 2O 2 solutions by several degrees. We have modified this experiment so that it can be done on a small-scale basis, eliminating many of these problems and giving highly reproducible results. The H2O2 stock solution is obtained from 3% supermarket or drug store hydrogen peroxide. Similar results are obtained from 30% laboratory H2O 2 that has been diluted to 3%. The apparatus for this experiment is shown in Figure 1. A 2-mL graduated Beral pipet is cut at each of the two arrows. A yellow micropipet tip is inserted into a number 2 one-hole stopper as shown in the completed assembly. The small bubbles released by the micropipet tip allow lower concentrations of hydrogen peroxide to be used. The barrel *Corresponding author.
of the pipet is fitted firmly over the tip and the stopper is placed in the 20-mL glass bottle reaction vessel. The water level is the same in each experiment so that the bubbles are pushing against the same pressure each time. (Too much pressure promotes the formation of double bubbles.) A preliminary check on the rate of decomposition of H2 O2 is made by using a 0.5–1.5-g piece of pyrolusite. Loose dust on the mineral is removed by swirling it in water with tweezers (it should not be handled with fingers) and blotting it dry on paper towels. Two milliliters of 3% H2O 2 is added to the reaction vessel and the catalyst is placed in the vial with tweezers. It is necessary for the parent stock solution to decompose at a rate close to but not greater than 2 b/s (bubbles of oxygen per second). It may be necessary to use a different sized piece of pyrolusite in order to have an appropriate decomposition rate. A detailed experimental procedure may be obtained from Ronald O. Ragsdale. Five concentrations of H2O2 are used to obtain the order of decomposition of the H2O 2. To maintain a constant temperature during the reaction, 10 mL of solution in the reaction bottle is placed in a 6-oz. Styrofoam cup, which is nesting in a 250-mL beaker. Water is added to the styrofoam cup so that its level is just above the level of the reaction solution. After the system has equilibrated, the catalyst is added and measurements are begun after about one minute. (The H2O 2 solution needs to soak into the catalyst.) Calculations The order of the decomposition of hydrogen peroxide is obtained by a graphical method. Rate (in bubbles per second) = k[H2O2]m Taking the log of both sides of the equation gives log Rate = m log [H2O2] + log k so that the slope (m) of a plot of log Rate vs log [H2 O2 ] is the order of the reaction. Students report the order to the nearest 0.1 unit. The effect of temperature on reaction rate is found by using an appropriate concentration of H2O2 and adding hot water, room-temperature water, cold water, and ice to the cup to achieve four different temperatures. By plotting ln Rate vs 1/T, the slope of the line corresponds to {Ea /R where R = 8.31 J/mol ? K. Students calculate the activation energy from their graph. When this experiment was done by a class of 240 pairs of students, the order was found to be 1.1 ± 0.2 and the activation energy was 35 ± 14 kJ. A typical plot of log R vs log [H2 O2] is shown in Figure 2. Owing to the complexity of the heterogeneous catalysis of hydrogen peroxide, it shouldn’t be surprising that
JChemEd.chem.wisc.edu • Vol. 75 No. 2 February 1998 • Journal of Chemical Education
215
Log Rate
In the Laboratory
m = slope = 1.24
Figure 1. A Beral pipet is shown with arrows indicating where it is to be cut. It is inserted over the micropipet tip into the stopper. When filled with water, it can be used to monitor the formation of oxygen bubbles.
the reaction order is not exactly one. Reproducible data have also been obtained with the minerals psilomelane, manganite, and groutite as catalysts. In addition, manganese metal having a coating of MnO2 can be used (2) and a rapid catalytic decomposition of hydrogen peroxide has been observed with a dry cell that has been cut open. Our results are especially gratifying considering that it is often difficult to obtain reproducible decomposition rate with heterogeneous catalysts (3). The decomposition of hydrogen peroxide appears to be related to its ability to function as both an oxidizing and reducing agent (4), an idea that is further supported by studies indicating that catalysis takes place by successive oxidation and reduction of manganese (5, 6). It has been reported that the catalytic activity of manganese is greatest when the oxidation state of manganese is between +3 and +4, and an oxidation–reduction cycle involving Mn2O 3 and MnO2 has been proposed (7). It is interesting to note that the formula of the manganese dioxide obtained by reduction of MnO4{ in neutral solution seldom corresponds exactly to the formula MnO2 (2).
216
Log [H2O2] Figure 2. A plot of the log Rate of decomposition of hydrogen peroxide versus log of concentration. The slope is the order of the reaction.
Acknowledgments Support by the University of Utah Park Fellowship to Ronald O. Ragsdale and by the Fund for Improvement of Postsecondary Education of the United States Department of Education for Ronald O. Ragsdale and Jan C. Vanderhooft is gratefully acknowledged. We wish to thank Brian Bennett for generating Figure 2 from some student data. Literature Cited 1. Slater, A. Chem 13 News; September 1993, p 17. 2. Schumb, W. C.; Satterfield, C. N.; Wentworth, R. L. Hydrogen Peroxide; Reinhold: New York, 1955; p 496. 3. Ibid., p 360. 4. Ibid., p 355. 5. Broughton, D. B.; Wentworth, R. L. J. Am. Chem. Soc. 1947, 69, 741. 6. Broughton, D. B.; Wentworth, R. L.; Laing, M. E. J. Am. Chem. Soc. 1947, 69, 744. 7. Mooi, J.; Selwood, P. W. J. Am. Chem. Soc. 1952, 74, 1750.
Journal of Chemical Education • Vol. 75 No. 2 February 1998 • JChemEd.chem.wisc.edu
