In the Laboratory edited by
filtrates & residues
James O. Schreck University of Northern Colorado Greeley, CO 80639
Kinetics Studies Using a Washing Bottle John Teggins and Chris Mahaffy Department of Physical Sciences, Auburn University at Montgomery, Montgomery, AL 36117 The kinetics of the decomposition of hydrogen peroxide using iodide ion as a catalyst has been commonly included in beginning high school and college chemistry courses. The usually accepted reaction mechanism results in a rate law that is first order with respect to both hydrogen peroxide and iodide ion in neutral solution.
Table 1. Reactants for Kinetic Runs and Rate Data at 21 °C
Volume (mL) 3% H2O2 0.1 M KI
H2O2 + I{ → H2O + IO{ (slow) { { H2O2 + IO → H2O + I + O2(g) (fast) {
Rate = k[H2O2]1[I ]1 The progress of the reaction is usually followed by collecting the oxygen over water using a gas buret (1). If the reactants are mixed in a beaker and are used to completely fill a bottom- or syphon-feeding washing bottle that can then be tightly sealed, the oxygen gas produced forces liquid out of the spout of the washing bottle. After a two- or three-minute initiation period, the rate at which solution is expelled from the spout of the washing bottle forms a convenient and surprisingly accurate method of following the rate of the reaction while avoiding complications involved in collecting and measuring gas volumes. All procedures can be easily performed from a seated position, making the experiment convenient for a student who requires a wheelchair. A suitable procedure is described below. Experimental Procedure SAFETY PRECAUTIONS: Safety glasses should be worn. Mix 100 mL each of a 0.1 M potassium iodide solution, 3% hydrogen peroxide solution, and deionized water in a beaker. Quickly fill a 250-mL washing bottle with the reaction mixture and firmly replace the cap. (A 250-mL Nalgene Unitary Washing Bottle from Fisher Scientific was found to be suitable: Catalog #03-409-16C.) Clamp the washing bottle in a beaker containing water at room temperature to maintain a nearly constant reaction temperature. After three minutes, the solution is expelled from the bottle at a steady rate. This solution is collected in a graduated cylinder. The total expelled volume is recorded at one-minute intervals. Determine the average reaction rate in terms of
H2O
50
100
150
100
100
100
200
100
0
100
50
150
100
200
0
8.8
0.52
0.50
17
1.00
1.00
35
2.06
2.00
0.50
0.50
1.94
2.00
8.5 33
a
Considering the reaction containing 100 mL of each reactant to have an initial rate = 1.00. Relative rates were obtained by dividing the average rates for the individual reactions by the value of 17 mL/min obtained in the second experiment. b Assuming that rate = k [H 2O2]1 [KI] 1.
milliliters of expelled solution per minute for five minutes after the rate has become approximately constant. Because only a small fraction of the hydrogen peroxide has decomposed in this time, this value can be considered to be proportional to the initial reaction rate. Repeat the experiment using different initial amounts of reactants as demonstrated in Table 1. The typical data in Table 1 were obtained under conditions where the temperature of the solutions did not increase more than 1 °C over the course of the reaction. All procedures employed could easily be repeated by a student under normal laboratory conditions in a first chemistry course. Nevertheless, observed relative rates for the reaction mixtures were in good agreement with theoretical values for a reaction that was first-order with respect to both iodide and peroxide concentrations. The procedure could readily be performed at different temperatures to obtain a value for the activation energy of the reaction. Literature Cited 1. Sienko, M. J.; Plane, R. A.; Marcus, S. T. Experimental Chemistry, 6th ed.; McGraw-Hill, 1984; p 108.
Hydrogen Peroxide
566
Expected Ave. Rate Rel. Ratea Rel. Rateb {1 {1 (mL/min) (mol L s ) (mol L{1 s{1)
Journal of Chemical Education • Vol. 74 No. 5 May 1997
