1. Chapter 12: Chemical Kinetics
Reaction Rates
Rate Laws: Differential and Integrated
Experimental Determination of the Rate Law
Integrated Rate Laws & Concentration/Time Problems
Collision Theory of Reactions
Reaction Mechanisms
Predicting Rate Laws
Temperature Dependence of the Rate Constant (Arrhenius)
Catalysis

2. As the reaction proceeds, it gets slower because the rate depends on [NO2].

3. Figure 12.1

4. Initial Rate (mole/L s)
Example 1:
Determine the rate law and the value of the rate
constant for the following reaction.
2 NO Br2  2 NOBr
Experiment
[NO]0
[Br2]0
Initial Rate (mole/L s) 1
1.1 x 10-5
1.2 x 10-5
0.37 2
1.9 x 10-5
1.3 x 10-5
0.69 3
3.5 x 10-5
1.2 4
4.0 x 10-5
3.0 x 10-5
3.4

5. Example 2: Determine the rate law and the value of the
rate constant for the following reaction in which
OH- is a catalyst.
OCl I-  OI Cl-
Experiment
[OCl-]0
(mole/L)
[I-]0
[OH-]0
Rate
(mole/L s) 1
.0040
.0020
1.00
4.8x10-4 2
5.0x10-4 3
2.4x10-4 4
0.50
4.6x10-4 5
0.25
9.4x10-4

6. Example 3:
The concentration of H2O2 was monitored over time for the following reaction at 25°C:
H2O2 (aq)  H2 (g) + O2 (g)
Find the rate law and the value of the rate constant for this reaction at 25°C.
Time (s)
[H2O2]
(mole/L)
0.000
1.000
0.910
0.780
0.590
0.370
0.220
0.130
0.082
0.050

7. Example 4: The decomposition of N2O5 to NO2 and O2 is
first order with a rate constant of
4.80 x 10-4 /s at 45°C.
If the initial concentration of N2O5 is 1.65 x 10-2 mole/L, what is the concentration after 825 s?
How long would it take for the concentration of N2O5 to decrease to 1.00 x 10-2 mole/L if its initial concentration wa that given in (a)?

8. Example 5: For the following decomposition reaction, AB  A + B
Rate = k[AB]2
k = 0.20 L/mole s
How long will it take for [AB] to reach one
third of its original value of 1.50 M? What
is [AB] after 10.0 seconds?

9. Zero Order Half Life t1/2 = [A]0 / 2k Note that this half life
decreases as the reaction
proceeds.
There is less reactant to
Consume and it runs at the
same rate so it takes less
time.
Time (s)
[A]
1.6
2.0
.80
3.0
.40
3.5
.20
3.75
.10
3.88
.050

10. First Order Half Life t1/2 = 0.693 / k
The first order half life is constant
for a particular reaction and
temperature.
As the initial concentration
decreases, the decrease in rate
is compensated for by less
reactant to consume.
Time (s)
[A]
1.6
24.0
.80
48.0
.40
72.0
.20
96.0
.10
120.0
.050

11. Second Order Half Life t1/2 = 1 / [A]0 k
Note that this half life increases
as the reaction proceeds.
The rate decreases so strongly
with concentration that it still
takes longer to use up half of
the reactant even though the
amount of reactant consumed is
less for each half life.
Time (s)
[A]
1.6
10.2
.80
30.6
.40
71.4
.20
153
.10

13. Example 6: A decomposition reaction has a rate
constant of yr-1.
What is the half life of this reaction?
How long does it take for the concentration of the reactant to reach 12.5% of its original value?

14. Example 7: It took 143 s for 50.0% of a particular
substance to decompose. If the initial
concentration was M and the
decomposition reaction follows second
order kinetics, what is the value for the
rate constant?

15. Example 8: For the reaction A  products,
successive half-lives are observed to be
10.0, 20.0, and 40.0 min for an
experiment in which [A]0=0.10M.
Calculate the concentration of
A at the following times.
80.0 min
30.0 min

16. Example 9: (Book #34) O + NO2  NO + O2 [NO2] in large excess,
1.0x1013molecules/cm3
Find the order of the reaction with respect to [O].
Reaction is known to be first order in [NO2]. Find the value of k.
Time (s)
[O] atoms/cm3
5.0x109
1.0x10-2
1.9x109
2.0x10-2
6.8x108
3.0x10-2
2.5x108

17. Example 10: Consider the following mechanism for the decomposition
of hydrogen peroxide catalyzed by I-.
Step 1: H2O2 + I H2O + IO- ; fast equilibrium
Step 2: IO- + H2O2  H2O + O2 + I- ; slow
a) Write a balanced equation for the overall reaction.
b) List any catalysts.
c) List any intermediates.
d) Write the rate law for the overall reaction.

18. Example 11: Consider the following mechanism for the production of
phosgene:
Step 1: Cl2 (g) Cl (g); fast equilibrium
Step 2: Cl (g) + CO (g) COCl (g); fast equilibrium
Step 3: COCl (g) + Cl2 (g)  COCl2 (g) + Cl (g); slow
Step 4: 2 Cl (g)  Cl2 (g); fast
a) Write a balanced equation for the overall reaction.
b) List any catalysts.
c) List any intermediates.
d) Write the rate law for the overall reaction.

19. Figure 12.11: 2BrNO  2 NO + Br2

20. Figure 12.13:

21. Example 12: The activation energy of the following
reaction is 3.5 kJ/mole and the change in
enthalpy is kJ/mole. Calculate
the activation energy of the reverse
reaction.
OH + HCl  H2O + Cl
Hint: Sketch the reaction energy diagram.

22. Example 13: For the reaction A2 + B2  2 AB, the
activation energy of the forward reaction is
125 kJ/mole and that of the reverse
reaction is 85 kJ/mole. Assuming the
reaction occurs in one step:
Draw a reaction energy diagram.
Calculate DH for the reaction.
Sketch a possible transition state.

23. Figure 12.2: Collisions with E > Ea at two temperatures.

24. Example 14: The isomerization reaction: CH3NC  CH3CN
has an activation energy of 161 kJ/mole.
If the rate constant at 600K is 0.41/s,
what is it at 1000K?

25. Example 15: The rate constant of a reaction is
4.50x10-5 L/mol•s at 195°C and
3.20 x 10-3 L/mol•s at 258°C.
What is the activation energy of this
reaction?

26. Catalysts Work by Lowering Ea: They provide a different mechanism!

27. More collisions result in reaction when Ea is lower.

28. Homogeneous Catalysis
In the stratosphere, NO catalyzes the
destruction of ozone, O3.
NO (g) + O3 (g)  NO2 (g) + O2 (g)
O (g) + NO2 (g)  NO (g) + O2 (g)
Overall:
O3 (g) + O (g)  2 O2 (g)

29. Heterogeneous Catalysis
C2H4 + H2  C2H6 with a metal catalyst a. reactants b. adsorption c. migration/reaction d. desorption

30. Enzymes are Catalysts

31. Enzymes are Catalysts

32. Enzymes are Catalysts