CHE1031 Lecture 10: Reaction kinetics Lecture 10 topics Brown chapter 14 1. Reaction rates Factors that effect reaction rates 14.1 Visualizing rates & units 14.2 Average reaction rates Instantaneous reaction rates Stoichiometry & reaction rates 2. Concentration & reaction rates 14.3 Rate laws Reaction orders 3. Change in concentration with time 14.4 First- & second-order reactions Half-life 4. Temperature & reaction rate 14.5 Collision, orientation & Ea 5. Reaction mechanisms 14.6 Elementary Multistep 6. Catalysis ` 14.7

Elementary vs. multistep reactions Reaction mechanisms Elementary vs. multistep reactions Rate-limiting steps The general process of advancing scientific knowledge by making experimental observations and by formulating hypotheses, theories, and laws. It’s a systematic problems solving process AND it’s hands-on….. Experiments must be done, data generated, conclusions made. This method is “iterative”; it requires looping back and starting over if needed. [Why do you think they call it REsearch?] Often years, decades or more of experiments are required to prove a theory. While it’s possible to prove a hypothesis wrong, it’s actually NOT possible to absolutely prove a hypothesis correct as the outcome may have had a cause that the scientist hasn’t considered.

Reaction mechanisms Reaction mechanism – describe the details by which a chemical rxn takes place. There are a series of basic & common mechanisms. Elementary mechanisms: single-step reactions Multistep reactions: require several steps, or reactions to reach completion Molecularity? Uni-, bi-, or termolecular describe reactions of 1, 2 or 3 molecules. Unimolecular: H3C – N = C  H3C – C = N = = Bimolecular: NO + O3  NO2 + O2 Two-step reaction: (1) NO2 + NO2  NO3 + NO (2) NO3 + CO  NO2 + CO2 sum NO2 + CO  NO + CO2 Molecules that don’t appear in the summed (overall) reaction are intermediates. p. 581 - 2

Elementary reactions & rate laws The relationship is quite simple as you can see here: p. 583 - 4

Rate-limiting steps When chemical reactions require more than one step, their overall rate is often limited by the slowest of the steps. So this slowest stop is called the rate-limiting step because it limits the overall rate of reaction. Step 1: NO2 + NO2  NO3 + NO (slow) Step 2: NO3 + CO  NO2 + CO2 (fast) Overall: NO2 + CO  NO + CO2 k1 k2 >> k1 k2 What is the rate law of the overall reaction? Because step 1 is much slower than step 2, it is rate-limiting. The rate of the overall reaction is equal to the rate of the slow step (1). Step 1 is bimolecular, so rate = k1[NO2]2 p. 584 - 5

Rate-limiting examples Nitrous oxide decomposes by a two-step mechanism. N2O  N2 + O (slow) N2O + O  N2 + O2 (fast) Write the equation for the overall reaction. Write the rate law for the overall reaction. a) 2N2O  2N2 + O2 b) Rate = k[N2O]2 Ozone reacts with nitrogen dioxide by a two-step mechanism: Step 1: O3 + NO2  NO3 + O2 Step 2: NO3 + NO2  N2O5 Overall: O3 + 2NO2  N2O5 + O2 Overall experimental rate law is: rate = k[O3][NO2]. Which step is slower? Step 1 is the slow step, since it is used in the rate law for the overall reaction. p. 584 - 5