Hydrogen peroxide decomposes slowly in neutral aqueous solution, but the reaction is greatly accelerated by iodide ion, which is recovered unchanged after the reaction.
Overall stoichiometry: 2 H2O2(aq) → 2 H2O(l) + O2(g)
A mechanism consistent with several experimental observations is proposed below (species are aqueous unless otherwise noted):
Step 1 I⁻ + H2O2 ⇌ HOI + OH⁻ fast equilibrium (forward constant k1, reverse constant k−1)
Step 2 HOI + H2O2 → I⁻ + H2O + O2 slow (rate constant k2)
Assume that water is the solvent and its concentration is effectively constant.
(a) Write the rate law for the slow elementary step in terms of the concentrations of HOI and H2O2.
(b) Using the equilibrium in Step 1, express [HOI] in terms of [I⁻], [H2O2], and the equilibrium constant Kc=k−1k1.
(c) Derive the overall rate law in terms of only [I⁻] and [H2O2].
(d) Determine the reaction order with respect to the catalyst I⁻. Explain how this order is reflected in the overall rate law.
(e) Predict quantitatively how the initial rate will change if the initial concentration of I⁻ is tripled while [H2O2] is unchanged.