1. Reaction Mechanisms
Happiness Factory
The chance of more than two particles colliding simultaneously with correct geometry and minimum energy required is very small.
If there are more than 2 reactants, the reaction will most likely occur by a number of simpler steps
A reaction mechanism is the step or series of steps that make up a reaction.

2. Reaction Mechanisms
A reaction mechanism is a series of simple steps that ultimately lead from the initial reactants to the final products of a reaction.
A valid mechanism must satisfy the following 2 criteria:
be consistent with the stoichiometry of the overall or net reaction.
account for the experimentally determined rate law.

3. Reaction Mechanisms
Molecularity : refers to the number of reactants involved in an elementary step. Usually limited to 1 or 2, rarely 3.
Consider the production of NO2 from NO and O2.
2 NO + O2  2 NO2
Is it likely for this reaction to proceed in one step? Explain
This reaction consists of 2 elementary steps.
1) 2 NO  N2O2 (fast)
2) N2O2 + O2  2 NO2 (slow, rate determining)
N2O2 is a reaction intermediate which is short lived and difficult to isolate.

4. Often it is difficult to determine the actual mechanism of an reaction, most are proposals supported by experimentation. (Rate Law)
The rate-determining step is the crucial step in establishing the rate of the overall reaction.
The rate determining step is the slowest step in the reaction mechanism (bottle neck). Must relate to the experimentally determined rate law equation.
The rate determining step will therefore have the highest activation energy.

5. Rate Law & Elementary Steps
Can use the equation coefficients as the reaction orders in the rate law for an elementary step

6. Rate Law Relationship
Lets revisit the the production of NO2 from NO and O2.
2 NO + O2  2 NO2
The proposed mechanism consists of 2 elementary steps.
1) 2 NO  N2O2 (fast)
2) N2O2 + O2  2 NO2 (slow, rate determining)
If this is a valid mechanism what is the rate law equation for this reaction?
Rate = k[O2]
The equation reflects the rate determining (slow) step.

7. Catalyzed Decomposition of Hydrogen Peroxide
Slow step: H2O2 + I-  H2O + OI-
Fast step: H2O2 + OI-  H2O + O2 + I-
___________________________________________________________
Net equation: 2 H2O2  2 H2O + O2
The slow step is the rate-determining step.
Rate of the reaction = rate of slow step = k[H2O2][I-]
OI- = ?
- Reaction intermediate
I- = ?
- Catalyst

8. Catalyzed Reactions Reaction Intermediate
A species that is created in one step and consumed in the other
Catalyst
A species that is present originally then reforms later on during the reaction
It is not written in the overall equation, but you may see it noted above the reaction arrow.
Catalysts provide a lower activation energy pathway by producing an intermediate.

9. Homework
Read section 6.6
Page 386 #1-3
Page 387 #3,5