1. Chemical Kinetics
The area of chemistry that examines reaction rates in order to understand the path of a reaction.
Thermodynamics is a state specific study of reactions. The path by which the products are formed is not considered. Kinetics is the study of the path of a chemical reaction. This path is called a mechanism.

2. Factors That Affect Reaction Rates
Physical State of the Reactants
In order to react, molecules must come in contact with each other.
The more homogeneous the mixture of reactants, the faster the molecules can react.

3. Factors That Affect Reaction Rates
Concentration of Reactants
As the concentration of reactants increases, so does the likelihood that reactant molecules will collide.

4. Factors That Affect Reaction Rates
Temperature
At higher temperatures, reactant molecules have more kinetic energy, move faster, and collide more often and with greater energy.

5. Factors That Affect Reaction Rates
Presence of a Catalyst
Catalysts speed up reactions by changing the mechanism of the reaction.
Catalysts are not consumed during the course of the reaction.

6. Path Discovery Process
Balanced equation
Consider possible paths
Run Rate Experiments
Determine Rate Law
Describe path which is consistent with the Rate Law
The process of investigation leading to a consistent path explanation is outlined above.

7. Reaction Rates
Rates of reactions can be determined by monitoring the change in concentration of either reactants or products as a function of time.

8. Reaction Rate
Change in concentration (conc) of a reactant or product per unit time.
The rate of a reaction can be examined by the disappearance of reactant or the appearance of product. Usually one is colored and can be quantified over time.

9. C4H9Cl(aq) + H2O(l)  C4H9OH(aq) + HCl(aq)
In this reaction, the concentration of butyl chloride, C4H9Cl, was measured at various times.

10. C4H9Cl(aq) + H2O(l)  C4H9OH(aq) + HCl(aq)
The average rate of the reaction over each interval is the change in concentration divided by the change in time:
Average rate =
[C4H9Cl] t

11. C4H9Cl(aq) + H2O(l)  C4H9OH(aq) + HCl(aq)
Note that the average rate decreases as the reaction proceeds.
This is because as the reaction goes forward, there are fewer collisions between reactant molecules.

12. C4H9Cl(aq) + H2O(l)  C4H9OH(aq) + HCl(aq)
A plot of concentration vs. time for this reaction yields a curve like this.
The slope of a line tangent to the curve at any point is the instantaneous rate at that time.

13. C4H9Cl(aq) + H2O(l)  C4H9OH(aq) + HCl(aq)
All reactions slow down over time.
Therefore, the best indicator of the rate of a reaction is the instantaneous rate near the beginning.

14. Reaction Rates and Stoichiometry
C4H9Cl(aq) + H2O(l)  C4H9OH(aq) + HCl(aq)
In this reaction, the ratio of C4H9Cl to C4H9OH is 1:1.
Thus, the rate of disappearance of C4H9Cl is the same as the rate of appearance of C4H9OH.
Rate =
-[C4H9Cl] t =
[C4H9OH]

15. Reaction Rates and Stoichiometry
What if the ratio is not 1:1?
2 HI(g)  H2(g) + I2(g)
Therefore,
Rate = − 1 2
[HI] t =
[I2]

16. Reaction Rates and Stoichiometry
To generalize, then, for the reaction
aA + bB
cC + dD
Rate = − 1 a
[A] t
= − b
[B] = c
[C] d
[D]

17. Rate Laws Rate = k[NO2]n k = rate constant n = rate order
After a reaction has been studied, the data is used to create a rate law. This law describes the reaction collisions in the slowest reaction in the controlled environment. The rate of the reaction will change with the concentration of reactants. How the rate changes determines the order. The order is what is determined during a kinetics experiment.

18. Types of Rate Laws
Differential Rate Law: expresses how rate depends on concentration. – Former Slide
Integrated Rate Law: expresses how concentration depends on time. – Examples to follow. (3 slides from now)
Definition Slide

19. Method of Initial Rates
Initial Rate: the “instantaneous rate” just after the reaction begins.
The initial rate is determined in several experiments using different initial concentrations.
The initial rate method is used because we know the intial concentrations. After the reaction has taken place, the concentrations have changed. We can make some assumptions in the math when using initial concentrations. Many times, we will only vary one reactant. This allows use to determine the order of that reactant. This is called pseudo order.

20. Overall Reaction Order
Sum of the order of each component in the rate law.
rate = k[H2SeO3][H+]2[I-]3
The overall reaction order is = 6.
Definition Slide

21. First-Order Rate Law For a reaction where
aA ® Products in a 1st-order reaction,
Look at the integrated rate law. What would one graph for a straight line? Also, if the reaction is half complete, what does time equal?
Integrated first-order rate law is
ln[A] = -kt + ln[A]o

22. Figure 12.4 A Plot of In(N2O5) Versus Time

23. Half-Life of a First-Order Reaction
t1/2 = half-life of the reaction
k = rate constant
For a first-order reaction, the half-life does not depend on concentration.

24. Figure 12.5 A Plot of (N2O5) Vs. Time for the Decomp. of N2O5

25. Second-Order Rate Law For aA ® products in a second-order reaction,
Integrated rate law is
Again, what is graphed for a straight line?

26. Figure 12.6 (a) A Plot of In(C4H6) Versus t (b) A Plot of 1/(C4H6) Versus t

27. Half-Life of a Second-Order Reaction
t1/2 = half-life of the reaction
k = rate constant
Ao = initial concentration of A
The half-life is dependent upon the initial concentration.

28. Figure 12.7 A Plot of (A) Versus t for a Zero-Order Reaction

29. A Summary
Simplification: Conditions are set such that only forward reaction is important.
Two types:
differential rate law
integrated rate law
Both forms can be interconverted.

30. A Summary (continued) Most common: method of initial rates.
Concentration v. time: used to determine integrated rate law, often graphically.
For several reactants: choose conditions under which only one reactant varies significantly (pseudo conditions).