Introduction

The decomposition of hydrogen peroxide in aqueous solution proceeds very slowly. A bottle of 3% hydrogen peroxide sitting on a grocery store shelf is stable for a long period of time. The decomposition takes place according to the reaction below.

{\text{2 }}{{\text{H}}_{\text{2}}}{{\text{O}}_{\text{2}}}\left( {{\text{aq}}} \right) \to {\text{2 }}{{\text{H}}_{\text{2}}}{\text{O }} + {\text{ }}{{\text{O}}_{\text{2}}}\left( {\text{g}} \right)

A number of catalysts can be used to speed up this reaction, including potassium iodide, manganese (IV) oxide, and the enzyme catalase. If you conduct the catalyzed decomposition of hydrogen peroxide in a closed vessel, you will be able to determine the reaction rate as a function of the pressure increase in the vessel that is caused by the production of oxygen gas. If you vary the initial molar concentration of the H2O2 solution, the rate law for the reaction can also be determined. Finally, by conducting the reaction at different temperatures, the activation energy, Ea, can be calculated.

Objectives

In this experiment, you will

  • Conduct the catalyzed decomposition of hydrogen peroxide under various conditions.
  • Calculate the rate constant for the reaction.
  • Determine the rate law expression for the reaction.
  • Calculate the activation energy for the reaction.