1. Reaction Mechanisms Elementary Steps Phenomenon Examples Rate Law
Unimolecular
Bimolecular
Termolecular
Rate-Limiting Step
How to Recognize
Reaction Intermediates
How to Recognize
Catalysts
How to Recognize

2. Elementary Reactions Main Concept:
Elementary reactions can be unimolecular or involve collisions between two or more molecules.

3. -Reactions occur when atoms and compounds collide with the right amount of energy and in the correct orientation.
-Often, multiple steps occur before a reaction is complete
-Elementary reactions are reactions that do not have multiple steps
ex:

4. Baking Soda and Vinegar
NaHCO3 + HC2H3O2 
H2O + CO2 + NaCH3COO
Actually, two steps…
NaC2H3O2 + H2CO3
H2CO3  H2O + CO2

5. Termolecular
Unimolecular
Bimolecular
- The order of an elementary reaction can be inferred from the number of molecules participating in a collision: unimolecular reactions are first order, reactions involving bimolecular collisions are second order, etc.

6. - Elementary reactions involving the simultaneous collision of three particles are rare
Termolecular

7. Multistep Reactions Main Concept:
The mechanism of a multistep reaction consists of a series of elementary reactions that add up to the overall reaction.

8. - The rate law of an elementary step is related to the number of reactants, as accounted for by collision theory.
Unimolecular A  Products rate = k[A]
Bimolecular A + A  Products rate = k[a]2
Bimolecular A + B  Products rate = k[A][B]
Termolecular A + A + A  Products rate = k[A]3
Termolecular A + A + B  Products rate = k[A]2[B]
Termolecular A + B + C  Prodcuts rate = k[A][B][C]

9. 2 NO2 (g) → NO3(g) + NO(g) NO3(g) + CO(g) → CO2 (g) + NO2(g)
- The elementary steps add to give the overall reaction
- The balanced chemical equation for the overall reaction shows only the stoichiometry of the reaction, not the rate.
Example:
This reaction…
CO(g) + NO2 (g) → CO2 (g) + NO(g)
is actually made up of two reactions…
2 NO2 (g) → NO3(g) + NO(g) NO3(g) + CO(g) → CO2 (g) + NO2(g)

10. - A number of mechanisms may be postulated for most reactions, and experimentally determining the dominant pathway of reactions is a central activity of chemistry.
*Each hump represents an elementary reaction

11. Rate Limiting Step Main Concept:
In many reactions, the rate is set up by the slowest elementary reaction, or rate-limiting step.

12. - For reactions where each elementary step is irreversible, the rate of reaction is set by slowest elementary step (ie. the rate-limiting or rate-determining step).

13. 2 NO2 (g) → NO3(g) + NO(g) NO3(g) + CO(g) → CO2 (g) + NO2(g)
Example:
Given the reaction…
CO(g) + NO2 (g) → CO2 (g) + NO(g)
which is actually made up of two reactions…
2 NO2 (g) → NO3(g) + NO(g) NO3(g) + CO(g) → CO2 (g) + NO2(g)
The rate law is rate = k[NO2]2
Because CO is not in the rate law (or doesn’t dramatically affect it), NO2 must be the slowest reaction and is called the rate determining step.

14. *The tallest hump (X3) represents the slowest elementary reaction

16. Reaction Intermediates
Main Concept:
Reaction intermediates, which are formed during the reaction but not present in the overall reaction, play an important role in multistep reactions.

17. - A reaction intermediate is produced by some elementary steps and consumed by others; it is present only while a reaction is occurring
- Experimental detection of a reaction intermediate is a way to build evidence in support of one reaction mechanism over an alternative mechanism
- Catalysts appear at the beginning of reaction steps and again at the end of the reaction steps, but not in the overall reaction

18. 2 NO2 (g) → NO3(g) + NO(g) NO3(g) + CO(g) → CO2 (g) + NO2(g)
Example:
Given the reaction…
CO(g) + NO2 (g) → CO2 (g) + NO(g)
which is actually made up of two reactions…
2 NO2 (g) → NO3(g) + NO(g) NO3(g) + CO(g) → CO2 (g) + NO2(g)
The reaction intermediate would include NO3
You can tell because it is not present at the beginning of the reaction nor the end

19. Example:
Given the reaction…
O3 + O  2 O2
which is actually made up of two reactions…
Cl + O3  ClO + O2
ClO + O  Cl + O2
The catalyst would be Cl (an intermediate would be ClO)
You can tell because it is present at the beginning of the reaction and the end, but not in the overall reaction