By Jack Brubaker, Updated Aug 30, 2022
Image credit: Apiwan Borrikonratchata/iStock/GettyImages
The iodine‑clock reaction is a classic demonstration used by high‑school and college chemistry students to visualize the principles of chemical kinetics. In this reaction, hydrogen peroxide oxidizes iodide to iodine. The iodine then reacts with thiosulfate until the thiosulfate is consumed. Once the thiosulfate is depleted, a starch indicator turns the solution a deep blue, marking the moment of the “clock.”
Every chemical transformation requires the breaking of bonds in the reactants. The energy that must be supplied to reach this transition state is known as the activation energy (Ea). While a reaction may be thermodynamically favorable—producing products with lower overall energy—the reaction rate is governed by Ea.
To determine Ea, one measures the rate constant (k) at several temperatures. Plotting the natural logarithm of k against the reciprocal of the absolute temperature (1/T, with T in Kelvin) should produce a straight line. The slope of this line equals –Ea/R, where R is the ideal gas constant (8.314 J mol⁻¹ K⁻¹).
For the iodine‑clock system, the ln k versus 1/T plot yields a slope of approximately –6 230. Using the relationship –Ea/R = –6 230 gives an activation energy of about 51.8 kJ mol⁻¹ (51 800 J mol⁻¹). This value reflects the energy barrier that must be overcome for the iodide oxidation and subsequent thiosulfate consumption to proceed.
