Reaction Mechanism

1. Reaction Mechanism The reaction mechanism is the series of elementary steps by which a chemical reaction occurs. • The sum of the elementary steps must give the overall balanced equation for the reaction • The mechanism must agree with the experimentally determined rate law

2. Rate-Determining Step In a multi-step reaction, the slowest step is the rate-determining step. It therefore determines the rate of the reaction. The experimental rate law must agree with the rate-determining step

3. Identifying the Rate-Determining Step For the reaction: 2H2(g) + 2NO(g)  N2(g) + 2H2O(g) The experimental rate law is: R = k[NO]2[H2] Which step in the reaction mechanism is the rate-determining (slowest) step? Step #1H2(g) + 2NO(g)  N2O(g) + H2O(g) Step #2N2O(g) + H2(g)  N2(g) + H2O(g) Step #1agrees with the experimental rate law

4. Identifying Intermediates For the reaction: 2H2(g) + 2NO(g)  N2(g) + 2H2O(g) Which species in the reaction mechanism are intermediates (do not show up in the final, balanced equation?) Step #1 H2(g) + 2NO(g)  N2O(g) + H2O(g) Step #2 N2O(g) + H2(g)  N2(g) + H2O(g) 2H2(g) + 2NO(g)  N2(g) + 2H2O(g)  N2O(g) is an intermediate

5. Identifying Catalysts For the reaction: O3(g) + O(g)  2O2 Intermediates are products in step 1 and reactants in step 2 Catalysts are reactants in step 1 and products in step 2 Step #1 O3(g) + NO(g)  NO2(g) + O2(g) Step #2 NO2(g) + O(g)  NO(g) + O2(g) O3(g) + O(g)  2O2  NO is the catalyst

6. Collision Model Key Idea: Molecules must collide to react. However, only a small fraction of collisions produces a reaction. Why?

7. Collision Model Collisions must have sufficient energyto produce the reaction (must equal or exceed the activation energy). Colliding particles must be correctly orientedto one another in order to produce a reaction.

8. Factors Affecting Rate Increasing temperature always increases the rate of a reaction. • Particles collide more frequently • Particles collide more energetically Increasing surface area increases the rate of a reaction Increasing Concentration USUALLY increases the rate of a reaction Presence of Catalysts, which lower the activation energy by providing alternate pathways

9. Arrhenius Equation k = Ae-Ea/RT , k is rate constant, A is frequency factor, e-Ea/RT is the fraction of collisions with sufficient energy to produce a reaction, Ea is activation energy, R is 8.31 J/mol K ln(k) = -Ea/R(1/T) + ln(A) linear equation ln(k2/k1) = Ea/R(1/T1 – 1/T2)

10. Catalysis • Catalyst: A substance that speeds up a reaction without being consumed • Enzyme: A large molecule (usually a protein) that catalyzes biological reactions. • Homogeneous catalyst: Present in the same phase as the reacting molecules. • Heterogeneous catalyst: Present in a different phase than the reacting molecules.

11. Lowering of Activation Energy by a Catalyst

12. Catalysts Increase the Number of Effective Collisions

13. Heterogeneous Catalysis Carbon monoxide and nitrogen monoxide adsorbed on a platinum surface Step #1: Adsorption and activation of the reactants.

14. Heterogeneous Catalysis Carbon monoxide and nitrogen monoxide arranged prior to reacting Step #2: Migration of the adsorbed reactants on the surface.

15. Heterogeneous Catalysis Carbon dioxide and nitrogen form from previous molecules Step #3: Reaction of the adsorbed substances.

16. Heterogeneous Catalysis Carbon dioxide and nitrogen gases escape (desorb) from the platinum surface Step #4: Escape, or desorption, of the products.