1. Kinetics Part V: Reaction Mechanisms
Jespersen Chap. 14 Sec 7
Dr. C. Yau
Fall 2013

2. Mechanisms
Mechanism refers to the series of individual steps that add up to the overall observed reaction. It tells us on the molecular level what is happening, such as what collides with what, and which the slowest step is.
Elementary reaction refers to each of the individual steps in the mechanism.
The sum of the elementary reactions must give the overall reaction.

3. Reaction Mechanisms
tell what happens on the molecular level, and in what order
tell us which steps in a reaction are fast and slow
The rate determining step is the slowest step of the reaction that accounts for most of the reaction time. 3

4. Example of a reaction mechanism
For the reaction of NO with H2
2NO + 2H  N H2O
The proposed mechanism is...
NO N2O (fast)
2. N2O2 + H2  N2O + H2O (slow)
3. N2O H2  N2 + H2O (fast)
These are the elementary rxns. Verify they add to give the overall equation.
Which is the rate determining step (rds)?

5. Elementary Steps: Molecularity vs. Rate Law
Elementary Reactions
Molecularity
Rate Law for the elementary step
A→Products
Unimolecular
Rate=k[A]
A+A →Products
A+B →Products
Bimolecular
Rate=k[A]2
Rate=k[A][B]
A+A+B →Products
3A →Products
A+B+C →Products
Termolecular
Rate=k[A]2[B]
Rate=k[A]3
Rate=k[A][B][C]
Molecularity refers to how many species are involved in the step. It is reflected in the overall order of reaction in the rate law.
Unlike the overall rate law, rate law of the elementary rxn uses the coefficients as the order of reaction.

6. Rate Laws And Mechanisms
The majority of the reaction time is taken by the rate determining step.
Substances that react in this step have the greatest effect on the reaction rate.
The observed rate law usually matches the rate law based on the rate determining step where the order of each reactant is its stoichiometric coefficient.
Note: Observed rate law is NOT based on the stoichiometric coefficients of the overall rxn.
Note: "Order" does not refer to which comes first, but to the molecularity: unimolecular? bimolecular? termolecular? 6

7. The “Transition State” in a Mechanism
Let us consider a simple reaction...
A + BC AB C
Reactants Products
One possible mechanism is to have this happen in one step.
Bond forming and bond breaking happening at the same time.
Reactants
Products
Transition State PE
Reaction Coordinate

8. The “Transition State” in a Mechanism
Bond forming and bond breaking happening at the same time.
Reactants
Products
Transition State
The transition state is unstable and cannot be isolated.
It is the highest point of the mechanism in the PE diagram.
Transition state = T.S. = activated complex

9. The “Intermediate” in a Mechanism
Let us consider the same reaction...
A + BC AB C
Reactants Products
We can have a mechanism where 2 steps are involved.
Step 1: B–C bond breaks to form B + C
Step 2: A forms bond with B. C B C B B A
Reactant
B-C
A is fully bonded to B.
Bond begins to
break.
C leaves.
B is alone, as an intermediate.
A begins to form bond to B.
Reactant T.S Intermediate T.S Product

10. B is the intermediate.
B-C bond begins
to break.
A-B bond begins to form. A C B
Reactants
A + B-C
Products
A-B + C
T.S.
Intermediate
An intermediate is formed during a reaction and consumed before the end of the reaction. It cannot be isolated.

11. Intermediates versus Catalysts
Because catalysts interact with the reactant, they will appear in the mechanism
Catalysts are consumed in an early step and are regenerated in a subsequent step
Intermediates are temporary products.
Intermediates are formed in an early step and consumed in a later step
Catalysts are there at beginning & at the end of the rxn.
Intermediates are not there at the beginning, nor at the end of the rxn.
Neither will show up in the overall eqn.

12. Intermediates in a mechanism
2NO N2O (fast)
N2O2 + H2  N2O + H2O (slow)
N2O H2  N2 + H2O (fast)
Are there any catalysts or intermediates in this mechanism?
Catalysts are ADDED, and they reappear at the end.
Intermediates are not reactants. They appear somewhere in the elementary steps, but are consumed before the end.
(No catalyst; N2O2 and N2O are intermediates.)

13. Example 1: 2H2O2 2H2O + O2 Observed rate law: Rate = k[H2O2][I-]
Why is I- not in the chemical equation?
The reaction mechanism that has been proposed for the decomposition of H2O2 is
H2O2 + I H2O + IO- (slow)
H2O2 + IO H2O + O2 + I- (fast)
The 1st step is slowest and therefore the rds.
The rate law of the rds is Rate = k[H2O2][I-] Why?
This matches the observed rate law.
Do the elementary steps shown add up to the overall reaction?
The rds is said to be "bimolecular.“ Why?
Catalysts? Intermediates?

14. Example 2
The reaction: A B → D + F was studied and the following mechanism was finally determined
A + B → D (fast)
B + B → E (slow)
E → F (very fast)
What is the expected rate law of the overall reaction?
rate=k[B]2 Check to see that elementary steps add up to give the overall reaction.

15. Example 3 For the reaction 2NO2Cl (g) 2 NO2 (g) + Cl2 (g)
the elementary steps have been proposed as…
(1) NO2Cl (g) NO2 (g) + Cl (g)
(2) NO2Cl (g) + Cl (g) NO2 (g) + Cl2 (g)
Step 1 is "unimolecular."
Step 2 is "bimolecular."
Do they add up to give the overall reaction?
Write the rate law for each of the elementary steps.
Step 1 must be rds.
Expt Rate Law: Rate = k[NO2Cl]

16. The slowest step is the rate determining step (rds).
The rds limits the rate of the overall reaction so its rate law represents the rate law for the overall reaction. Example 4:
NO2 (g) + CO (g) NO (g) + CO2 (g)
The experimental rate law is
Rate = k[NO2]2
From this we know immediately the reaction shown cannot be elementary. Can you see why not?

17. Example 4 (cont'd.) NO2 (g) + CO (g) NO (g) + CO2 (g)
Proposed 2-step mechanism:
(1) NO2 + NO NO3 + NO
(2) NO3 + CO NO2 + CO2
Write the rate laws for these 2 elementary steps.
Given: experimental rate law: Rate = k[NO2]2
Which is the rate determining step?
How do you tell?

18. Example 5 2NO + 2H2 N2 + 2H2O Expt rate law: Rate = k[NO]2[H2]
which suggests 2NO + H N2O + H2O
and then N2O + H N2 + H2O
The 1st step proposed is termolecular, which is highly unlikely. Why?
It would require 3 molecules to all collide at the same time.

19. Example 5: 2NO + 2H2 N2 + 2H2O cont'd
Instead: 2NO N2O fast
N2O2 + H N2O + H2O slow
N2O + H N2 + H2O fast
If we look only at the rds (Step 2)
Rate = k[N2O2][H2] but N2O2 is not in the overall reaction, and not in the observed rate law.
This is because N2O2 is an intermediate.
When this happens, we have to find an expression to replace [N2O2] in the rate law.

20. 2NO N2O fast
N2O2 + H N2O + H2O slow
N2O + H N2 + H2O fast
For the equilibrium, forward reaction has…
Rate = kf [NO]2
and the reverse reaction has…
Rate = kr [N2O2]
At equilibrium, the two rates are equal, so
kf [NO]2 =kr [N2O2] which rearranges to
[N2O2] =?
Rate law for Step 2 is Rate = k[N2O2][H2]
Substitute [N2O2] from the equilibrium step and we have a rate law that matches the observed.

21. 2NO + 2H2 N2 + 2H2O Substitute [N2O2] = kf/kr [NO]2
into Rate = k [N2O2] [H2]
to give Rate = k kf/kr [NO]2[H2]
Since k, kf, kr are all constants, we can combine them into a new constant, k' to give
Rate = k'[NO]2[H2] and this matches the observed
Rate = k[NO]2[H2]
Do Practice Exercises
29, 30, 31 p.680

22. Potential Energy Diagram
One-step mechanism:
Potential E
1) Where is Ea?
2) Where is ΔH?
3) Is this exothermic or endothermic?
4) Where is the transition state?
5) What is the activated complex?
Reaction Coordinate

23. Potential Energy Diagram
Two-step mechanism:
8) Is this exo- or endothermic?
1) How do you know it has 2-steps?
2) What is at the "valley"?
3) Which step is rds?
4) Where is the ΔH of the rds?
5) How many transition states are there?
6) Which is the rds of the reverse reaction?
7) Where is the Ea of the 2nd step?
Potential E
Reaction Coordinate

24. Potential Energy Diagram
Two-step mechanism:
Potential E
1) In the reverse rxn, which is the rate determining step?
2) Where is the ΔH of the reverse rxn?
3) Is the reverse rxn exo- or endothermic?
Reaction Coordinate

25. Potential Energy Diagram
What can you say about this mechanism?
Potential E
How many steps are in the mechanism?
Is it exothermic or endothermic?
How many transition states are there?
What is at each "valley"?
Which step is rds?
Reaction Coordinate

26. PE Diagram of Reversible vs. Irreversible Rxns
Diagram A
Diagram B
Which PE diagram more closely corresponds to a reversible reaction? Why?

27. Potential Energy Diagram
What would the PE graph look like if a catalyst is added?
Graph for uncatalyzed rxn
Graph for catalyzed rxn
Potential E
Note that the Ea is lowered.
How does that affect the Kinetic E Diagram?
Reaction Coordinate

28. Kinetic Energy Diagram: Effect of Catalyst
Be able to explain why lowering the Ea in the KE diagram would affect the rate of reaction.

29. Applications of Catalytic Rxns
Oil industries: Catalysts help "crack" larger hydrocarbons into smaller pieces and reform them to more useful molecules.
Catalytic converter in your car: helps remove CO (poisonous), unburned volatile hydrocarbons and nitrogen oxides (contribute to acid rain)
Enzymes are biological catalysts: Lactase catalyzes the breakdown of lactose (Those who are lactose-intolerant lack lactase.)