Many advanced high school and college chemistry students perform an experiment known as the “iodine-clock” reaction, in which hydrogen peroxide reacts with iodide to form iodine, and the iodine subsequently reacts with thiosulfate ion until the thiosulfate has been consumed. At that point, the reaction solutions turn blue in the presence of starch. The experiment helps students understand the fundamentals of chemical kinetics — the speeds at which reactions take place.
Chemical reactions are thermodynamically “favorable” if the overall energy the the products is lower than the overall energy of the reactants. The formation of products, however, first requires bond breakage in the reactants, and the energy required to break them represents an energy barrier known as the “activation energy,” or Ea.
Measuring Activation Energy
The determination of activation energy requires kinetic data, i.e., the rate constant, k, of the reaction determined at a variety of temperatures. The student then constructs a graph of ln k on the y-axis and 1/T on the x-axis, where T is the temperature in Kelvin. The data points should fall along a straight line, the slope of which is equal to (-Ea/R), where R is the ideal gas constant.
Iodine-Clock Activation Energy
The plot of (ln k) vs. (1/T) for the iodine clock reaction should reveal a slope of about -6230. Thus, (-Ea/R) = -6230. Using an ideal gas constant of R = 8.314 J/K.mol gives Ea = 6800 * 8.314 = 51,800 J/mol, or 51.8 kJ/mol.