Reaction Mechanism See lectures 32 and 33 for determining reaction mechanisms. Example from pg. 4 of Lecture 33 notes: The reaction mechanism of natural ozone depletion: Through a variety of catalytic cycles, Cl and Br ions catalyze the destruction of ozone in the stratosphere. However, at high concentrations, ozone decays into oxygen without metal catalysts required: 2O3 → 3O2 proposed mechanism: k1 Step 1 O3 O2 + O (fast k-1 reversible) k2 Step 2 O + O3 → O2 + O2 (slow) ratef= k1[O3] rater= k-1[O2][O] rate= k2[O][O3] The rate is determined by the slowest step The rate of formation of O2 is equal to 2 times the rate of the slow step (k2[O][O3]), since two molecules of O2 are formed. Thus, rate of formation of O2 = 2k2[O][O3], but “O” is an intermediate, solve for “O” in terms of products and reactants and rate constants. Since the first step is fast and reversible and the second step is slow, the first step is in equilibrium and we can write [O2][O] = k1 = K1 [O3] k-1 or [O] = k1 [O3 ] k-1 [O2] substituting: rate = 2k2 k1 [O3 ]2 k-1 [O2] rate = kobs [O3]2 [O2] What is the order in O3? 2 What is the order in O2? -1 What is the overall order? 1 double O3/rate will? multiply by 4 double O2/ multiply by ½ double both O3 and O2/ double