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 #1 H2(g) + 2NO(g)  N2O(g) + H2O(g)
Step #2 N2O(g) + H2(g)  N2(g) + H2O(g)
Step #1 agrees 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 energy to produce the reaction (must equal or exceed the activation energy).
Colliding particles must be correctly oriented to 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.