1. CHAPTER 12
Chemical Kinetics

2. Kinetics The study of reaction rates.
Spontaneous reactions are reactions that will happen - but we can’t tell how fast.
Diamond will spontaneously turn to graphite – eventually.
Reaction mechanism- the steps by which a reaction takes place.

3. Reaction Rate Reaction Rate
Rate = Conc. of A at t2 -Conc. of A at t1 t2- t1
Rate = D[A] Dt
Change in concentration per unit time
For this reaction
N2 + 3H NH3

4. REACTION RATE GRAPH
As Time Increases, Concentration deceases

5. H2 VS N2
As the reaction progresses the concentration N2 goes down 1/3 as fast

6. PRODUCT CONCENTRATION
As the reaction progresses the concentration NH3 goes up.

7. Calculating Rates Average rates are taken over long intervals
Instantaneous rates are determined by finding the slope of a line tangent to the curve at any given point because the rate can change over time
Derivative.

8. Average slope method
SLOPE

9. Instantaneous slope method

10. Defining Rate
We can define rate in terms of the disappearance of the reactant or in terms of the rate of appearance of the product.
In our example
N2 + 3H2  NH3
-D[N2] = -3D[H2] = 2D[NH3] Dt Dt Dt

11. Rate Laws Reactions are reversible.
As products accumulate they can begin to turn back into reactants.
Early on the rate will depend on only the amount of reactants present.
We want to measure the reactants as soon as they are mixed.
This is called the Initial rate method.

12. Rate Laws Two key points
The concentration of the products do not appear in the rate law because this is an initial rate.
The order must be determined experimentally,
can’t be obtained from the equation

13. 2 NO2  2 NO + O2
You will find that the rate will only depend on the concentration of the reactants.
Rate = k[NO2]n
This is called a rate law expression.
k is called the rate constant.
n is the order of the reactant -usually a positive integer.

14. 2 NO2  2 NO + O2 The rate of appearance of O2 can be said to be.
Rate' = D[O2] = k'[NO2] Dt
Because there are 2 NO2 for each O2
Rate = 2 x Rate'
So k[NO2]n = 2 x k'[NO2]n
So k = 2 x k'

15. Types of Rate Laws
Differential Rate law - describes how rate depends on concentration.
Integrated Rate Law - Describes how concentration depends on time.
For each type of differential rate law there is an integrated rate law and vice versa.
Rate laws can help us better understand reaction mechanisms.

16. Determining Rate Laws
The first step is to determine the form of the rate law (especially its order).
Must be determined from experimental data.
For this reaction
2 N2O5 (aq)  4NO2 (aq) + O2(g) The reverse reaction won’t play a role

17. Now graph the data [N2O5] (mol/L) Time (s) 1.00 0 0.88 200 0.78 400

18. To find rate
To find rate we have to find the slope at two points We will use the tangent method.

19. At .90 M the rate is (0.98 M - .76 M) = 0.22 = -5.5 x10-4
(0 – 400)

20. At 0.40M the rate is: (.52 - .31) = 0.22 = -2.7 x 10-4 (1000-800) -800

21. WHAT IT MEANS
Since the rate at twice the concentration is twice as fast the rate law must be..
Rate = -D[N2O5] = k[N2O5]1 = k[N2O5] Dt
We say this reaction is first order in N2O5
The only way to determine order is to run the experiment.

22. The method of Initial Rates
This method requires that a reaction be run several times.
The initial concentrations of the reactants are varied.
The reaction rate is measured bust after the reactants are mixed.
Eliminates the effect of the reverse reaction.

23. An example For the reaction BrO3- + 5 Br- + 6H+ 3Br2 + 3 H2O
The general form of the Rate Law is Rate = k[BrO3-]n[Br-]m[H+]p
We use experimental data to determine the values of n,m,and p

24. Initial concentrations (M)
Rate (M/s)
BrO3- Br- H+
x 10-4
x 10-3
x 10-3
x 10-3
Now we have to see how the rate changes with concentration