1. Chapter 12
Chemical Kinetics

2. Kinetics Study of Speed at which reactions take place
Effect of each reactant
Effect of concentration
Mechanisms
Reaction process
Energy requirements for reactions

3. Reaction Rate
Change in concentration of reactant or product per unit of time
Rate = Δ[A]/ Δt
Rates decrease as time increases
Rates are given positive values
If you know the rate of one species in a reaction you can calculate the others

4. 2NO2  2NO + O2 Page 558 Based on the graph What do you notice?
Calculate the rate of
NO2 0s to 50.s
NO2 50.s to 100.s
O2 0s to 50.s
2NO2  2NO + O2

5. Rate Laws Only depends on the forward reaction
An expression that shows how the rate depends on the concentrations of reactants
A way of determining rate or concentration
Depends on what you know.

6. Form of the Rate Law Assuming the reaction A  B + C Rate = k[A]n
k and n are both experimentally determined
k is a constant called the rate constant
Units vary depending on n
n is the order
Does NOT depend on balanced equation
Whole numbers including zero and fractions

7. Units of k Depends on what number n is
k has units that gives the rate units of mol/L*s
Consider - Rate = k[A]n
What are the units of k if n is 1? 2? 0?

8. Multiple Reactants Consider the reaction A + B  C
What is the rate law?
Rate = k[A]n[B]m
What are the units for k when n is 1 & m is 1

9. Specifically These types of rate laws are differential rate laws
Just called rate laws
There are also integrated rate laws
The form you use depends on the data you have

10. Homework
P. 598 #’s 17,18,19,20

11. Method of Initial Rates

12. Integrated Rate Laws
A way of determining the order and rate constant of a reaction when time and concentration data is known.
Must be for a single reactant.
Or have one in excess
Equations for 0th, 1st, and 2nd order
Use equations as tests for order

13. 1st Order ln[A] = -kt + ln[A]o [A] concentration of reactant at time t
[A]o concentration of reactant at time t=0
k is the rate constant
This is the equation of a ________?
LINE
Y axis is ln[A], X axis is t
The slope is the rate constant

14. 1st Order
If a graph of ln[A] vs t is a straight line then the reaction is 1st order
A straight line is the check for all of these

15. 2N2O5 4NO2 + O2 Example Use the data to determine the rate law.
Determine the rate constant.
Determine the [N2O5] at 2000.s
Time (s)
[N2O5]
.100
100.
.0614
300.
.0233
600.
.00541
900.
.00126

16. 1st Order Half Life
The half life is the time required for a reactant to reach half of its previous concentration
Meaning [A] = [A]o/2
Derive Half Life Equation
t1/2=ln2/k
Half life is always the same for first order kinetics

18. 2N2O5 4NO2 + O2 Example Use the data to determine the rate law.
Determine the rate constant.
Determine the [N2O5] at 2000.s
Determine the half life
Time (s)
[N2O5]
.100
100.
.0614
300.
.0233
600.
.00541
900.
.00126

19. 2nd Order Integrated Rate Law 1/[A] = kt + 1/[A]o
This is the equation of a ________?
LINE
Y axis is 1/[A], X axis is t
The slope is the rate constant
Half Life
t1/2 = 1/(k[A]o)
Each half life is double the previous

20. 2C4H6  C8H8 Example #33 p. 600 Determine the rate law
Determine the rate constant.
Determine the half-life
Time (s)
[C4H8]
.017
195
.016
604
.015
1246
.013
2180
.011
6210
.0068

22. 0th Order Integrated Rate Law [A] = -kt + [A]o
This is the equation of a ________?
LINE
Y axis is [A], X axis is t
The slope is the rate constant
Half Life
t1/2 = [A]o/2k
Each half life is the same

23. 2N2O  2N2 + O2 Determine the rate law. Determine the rate constant.
Determine the half-life
Time (s)
[C4H8]
.44
10.
.33
20.
.22
30.
.11
40.

24. Page 578

25. Homework
Page 599 #’s 27,29,30,32

26. Reaction Mechanism
A series of elementary steps that must satisfy two requirements
The sum of the steps must equal the overall balanced equation
Must agree with the rate law
Elementary Step – Steps in the mechanism. (Think back to organic)

27. Cont.
Intermediates – Species that are produced in one elementary step and consumed in another
Not part of the overall balanced equation
Rate-determining step – Slowest step in a reaction that determines the rate
Molecularity – Number of species that must collide to produce an elementary step

28. Molecularity Table Elementary Step Molecularity Rate Law A  Products
Unimolecular
Rate=k[A]
2 A  Products
Bimolecular
Rate=k[A]2
A + B  Products
Rate=k[A][B]2
2A + B  Products
Termolecular
Rate=k[A]2[B]
A+B+C Products
Rate=k[A][B][C]

29. NO2 + CO  NO + CO2
The reaction above has an experimentally determined rate law of
Rate=k[NO2]2
Is the proposed mechanism possible? Explain
Step 1 = NO2 + NO2  NO3 + NO
Step 2 = NO3 + CO  NO2 + CO2

30. Collision Model for Kinetics
For a reaction to occur particles must collide to react.
Must collide in the proper orientation
Must collide with enough energy
Activation Energy

31. Proper Orientation + 
Images from:

32. Improper Orientation + 
Images from:

33. Activation Energy (Ea)
The energy required to convert atoms or molecules into their transition state
Minimum energy required for effective collisions
Can be found by running the same reaction at different temps.
Use Arrhenius Equation

34. Arrhenius Equation K = rate constant Ea = Activation Energy
R = Energy Gas Constant J/mol*K
T = Temp in K
A = Frequency Factor
Equation of a line Y is ln k X is 1/T
Slope is –Ea/R

35. Find the activation energy
2N2O5  4NO2 + O2
Rate Constant
Temp (ºC)
2.0x10-5 20
7.3x10-5 30
2.7x10-4 40
9.1x10-4 50
2.9x10-3 60

36. 2 Temp. Ea Equation
You can manipulate the Arrhenius Equ if you only have to temps and rate constants by subtracting the lower temp equation from the higher temp equation

37. Catalysts
Chemical that speeds up a chemical reaction but is not consumed in the reaction
Do this by
Lowering activation energy
Provide alternate reaction pathways
Increase effective collisions

38. Cont.
Below are two steps in the destruction of ozone. What is the overall rxn? What is the catalyst? What is the intermediate?
Cl + O3 ClO + O2
O + ClO  Cl + O2

39. Homework
Page 602 #’s 45,48,49,54