1. Why do some reactions happen and others don’t?
Are the products more stable than the reactants? Thermodynamics
Does the reaction go at a reasonable rate? Kinetics
Chapter 14: Chemical Kinetics Rates of Reactions

3. Control of Reactivity

4. Collision Theory Consider: NO + O3  NO2 + O2
For a reaction to take place:
- Molecules must collide
- They must do so in the correct orientation
- They must collide with an energy greater than the “activation” energy
Consider: NO + O3  NO2 + O2
product molecules
separate
Molecules collide
Bonds are
formed and
break

5. What would control how fast a reaction happens?

6. So, what controls the rate of a reaction?
Number of collisions
How often they collide in a shape that allows new bonds to form
The energy of the colliding reactant molecules
We’ll consider dependence on:
Concentration
a. Rate laws
b. Concentration vs. time relationships
Temperature and activation energy
Mechanism

7. Concentration Dependence
It makes sense that as concentration increases, the number of collisions per second will increase
Therefore, in general, as concentration increases, rate increases
But, it depends on which collisions control the rate
So, you can’t predict concentration dependence- it must be measured experimentally

8. Types of measured rates:
Rate over time:
Instantaneous rate:
Initial rate:

9. Example of rate measurement:

10. Rate Laws (also called Rate Equations)
first order reaction
For the reaction: 2 N2O5  4 NO + O2
Rate = k[N2O5]
For the reaction: NO2  NO + ½ O2
Rate = k[NO2]2
For the reaction: CO + NO2  CO2 + NO
Rate = k[CO][NO2]
second order reaction
first order in CO and in NO2; second order overall

11. Determining a Rate Law Method of Initial Rates
Determining the rate law must be done by experiment; the reaction equation does not tell you the rate law
Two methods: Initial Rates and the Graphical Method
Method of Initial Rates
Measure the rate of the reaction right at the start.
Vary the starting concentrations
Compare initial rates to initial concentrations

12. Determining a Rate Law: Initial Rate Method
Isolation of variables: Vary only one concentration at a time and keep temperature constant
If concentration doubles and:
Rate does not change, then zero order
Rate doubles, then first order
Rate quadruples, then second order
General Rule:

13. Initial Rate Method: Example 1
What is the rate law?

14. Initial Rate Method: Example 2

15. Concentration-Time Relationships

16. Example 1

17. Example 2 The decomposition of nitrous oxide at 565 oC,
2 N2O  2 N2 + O2
is second order in N2O. If the reaction is initiated with [N2O] equal to M,
and drops to M after 1250 s have elapsed, what is the rate constant?

18. Graphical Method for Determining Rate Laws
A plot of concentration
vs. Time will be linear.
A plot of ln[R]
vs. Time will be linear.
A plot of 1/[R]
vs. Time will be linear.

19. Graphical Method for Determining Rate Laws
How it works:
1. Collect [R] over an interval of times.
2. Make plots of
[R] vs. time
ln[R] vs. time
1/R vs. time
Only one will be linear. That tells you the reaction order.
The slope of the linear plot is the rate constant.

20. Graphical Method for Determining Rate Laws
Example: 2 H2O2  2 H2O + O2
Time(min) [H2O2](mol/L)

21. Graphical Method for Determining Rate Laws: Order
Example: 2 H2O2  2 H2O + O2
Time(min) [H2O2](mol/L)

22. Graphical Method for Determining Rate Laws: k

23. Half-Life: t1/2
the time it takes for half the reactant concentration to drop to half of its original value
First Order Reaction: 2 H2O2  2 H2O + O2 Rate = k[H2O2]; k = 1.05 x 10-3/min

24. Cool things about half-life:

25. Calculations involving Half-Life
For a first order reaction:
What is the relationship between t1/2 and k?
What is the relationship between t1/2 and k for a second order reaction?

26. Radioactive Decay
All radioisotopes decay via first order reactions. Instead of concentrations, amounts are used.
Measured as radioactive activity, in counts per minute (cpm) using a detector.

27. Radioactive Decay: Example 1
Radioactive gold-198 is used in the diagnosis of liver problems. The half-life of this isotope is 2.7 days. If you begin with a 5.6-mg sample of the isotope, how much of this sample remains after 1.0 day?

28. Radioactive Decay: Carbon Dating
In living thing
Sunlight + Nitrogen
Atmospheric C-14
C-14
Dead thing
Sunlight + Nitrogen
Atmospheric C-14

29. Radioactive Decay: Example 2
The Carbon-14 activity of an artifact in a burial site is found to be 8.6 counts per minute per gram. Living material has an activity of 12.3 counts per minute per gram. How long ago did the artifact die? t1/2 = 5730 years