1. CHAPTER 15
Chemical Kinetics

2. KINETICS
— the study of REACTION RATES and their relation to the way the reaction proceeds, i.e., its MECHANISM.
Reaction rate = change in concentration of a reactant or product with time.
Three “types” of rates
initial rate
average rate
instantaneous rate

3. The Rate of Reaction Consider the hypothetical reaction,
A(g) + B(g)  C(g) + D(g)
equimolar amounts of reactants, A and B, will be consumed while products, C and D, will be formed as indicated in this graph:

4. aA(g) + bB(g)  cC(g) + dD(g)
The Rate of Reaction
Mathematically, the rate of a reaction can be written as:
aA(g) + bB(g)  cC(g) + dD(g)

5. The Rate of Reaction
The rate of a simple one-step reaction is directly proportional to the concentration of the reacting substance.
[A] is the concentration of A in molarity or moles/L.
k is the specific rate constant.
k is an important quantity in this chapter.

6. The Rate of Reaction Important terminology for kinetics.
The order of a reaction can be expressed in terms of either:
each reactant in the reaction or
the overall reaction.
Order for the overall reaction is the sum of the orders for each reactant in the reaction.
For example:

7. In general, for
a A + b B --> x X with a catalyst C
Rate = k [A]m[B]n[C]p
The exponents m, n, and p
• are the reaction order
• can be 0, 1, 2 or fractions
• must be determined by experiment!

8. The Rate of Reaction
Given the following one step reaction and its rate-law expression.
Remember, the rate expression would have to be experimentally determined.
Because it is a second order rate-law expression:
If the [A] is doubled the rate of the reaction will increase by a factor of = 4
If the [A] is halved the rate of the reaction will decrease by a factor of 4. (1/2)2 = 1/4

9. Factors That Affect Reaction Rates
There are several factors that can influence the rate of a reaction:
The nature of the reactants.
The concentration of the reactants.
The temperature of the reaction.
The presence of a catalyst.
We will look at each factor individually.

10. Nature of Reactants
This is a very broad category that encompasses the different reacting properties of substances.
For example sodium reacts with water explosively at room temperature to liberate hydrogen and form sodium hydroxide.

11. Nature of Reactants
Calcium reacts with water only slowly at room temperature to liberate hydrogen and form calcium hydroxide.

12. Nature of Reactants
The reaction of magnesium with water at room temperature is so slow that that the evolution of hydrogen is not perceptible to the human eye.

13. Nature of Reactants
However, Mg reacts with steam rapidly to liberate H2 and form magnesium oxide.
The differences in the rate of these three reactions can be attributed to the changing “nature of the reactants”.

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15. Concentrations of Reactants: The Rate-Law Expression
This is a simplified representation of the effect of different numbers of molecules in the same volume.
The increase in the molecule numbers is indicative of an increase in concentration.
A(g) + B (g)  Products
A B B
A B
A B
A B
6 different possible A-B collisions
9 different possible A-B collisions
4 different possible A-B collisions

16. Concentrations of Reactants: The Rate-Law Expression
Example 16-1: The following rate data were obtained at 25oC for the following reaction. What are the rate-law expression and the specific rate-constant for this reaction? A(g) + B(g) ® 3 C(g)
Experiment
Number
Initial [A]
(M)
Initial [B]
Initial rate of formation of C (M/s) 1
0.10
2.0 x 10-4 2
0.20
0.30
4.0 x 10-4 3

17. Concentrations of Reactants: The Rate-Law Expression

18. Concentrations of Reactants: The Rate-Law Expression
The following data were obtained for the following reaction at 25oC. What are the rate-law expression and the specific rate constant for the reaction?
2 A(g) + B(g) + 2 C(g) ® 3 D(g) + 2 E(g)
Experiment
Initial [A]
(M)
Initial [B]
Initial [C]
Initial rate of formation of D (M/s) 1
0.20
0.10
2.0 x 10-4 2
0.30
6.0 x 10-4 3 4
0.60
0.40
1.8 x 10-3

19. Concentrations of Reactants: The Rate-Law Expression

20. Concentrations of Reactants: The Rate-Law Expression
Consider a chemical reaction between compounds A and B that is first order with respect to A, first order with respect to B, and second order overall. From the information given below, fill in the blanks.
Experiment
Initial Rate
(M/s)
Initial [A]
(M)
Initial [B] 1
4.0 x 10-3
0.20
0.050 2
1.6 x 10-2 ? 3
3.2 x 10-2
0.40

21. Concentrations of Reactants: The Rate-Law Expression

22. Concentration vs. Time: The Integrated Rate Equation
The integrated rate equation relates time and concentration for chemical and nuclear reactions.
From the integrated rate equation we can predict the amount of product that is produced in a given amount of time.
Initially we will look at the integrated rate equation for first order reactions.
These reactions are 1st order in the reactant and 1st order overall.

23. Concentration vs. Time: The Integrated Rate Equation
An example of a reaction that is 1st order in the reactant and 1st order overall is:
a A ® products
This is a common reaction type for many chemical reactions and all simple radioactive decays.
Two examples of this type are:
2 N2O5(g)  2 N2O4(g) + O2(g)
238U  234Th + 4He

24. Concentration vs. Time: The Integrated Rate Equation
The integrated rate equation for first order reactions is:
where:
[A]0= mol/L of A at time t=0. [A] = mol/L of A at time t.
k = specific rate constant. t = time elapsed since beginning of reaction.
a = stoichiometric coefficient of A in balanced overall equation.

25. Concentration vs. Time: The Integrated Rate Equation
Solve the first order integrated rate equation for t.
Define the half-life, t1/2, of a reactant as the time required for half of the reactant to be consumed, or the time at which [A]=1/2[A]0.

26. Concentration vs. Time: The Integrated Rate Equation
At time t = t1/2, the expression becomes:

27. Concentration vs. Time: The Integrated Rate Equation
Cyclopropane, an anesthetic, decomposes to propene according to the following equation.
The reaction is first order in cyclopropane with k = 9.2 s-1 at 10000C. Calculate the half life of cyclopropane at 10000C.

28. Concentration vs. Time: The Integrated Rate Equation
Refer to Previous Example: How much of a 3.0 g sample of cyclopropane remains after 0.50 seconds?
The integrated rate laws can be used for any unit that represents moles or concentration.
In this example we will use grams rather than mol/L.

29. Concentration vs. Time: The Integrated Rate Equation
The half-life for the following first order reaction is 688 hours at 10000C. Calculate the specific rate constant, k, at 10000C and the amount of a 3.0 g sample of CS2 that remains after 48 hours.
CS2(g) ® CS(g) + S(g)

30. Concentration vs. Time: The Integrated Rate Equation
For reactions that are second order with respect to a particular reactant and second order overall, the rate equation is:
Where:
[A]0= mol/L of A at time t=0. [A] = mol/L of A at time t.
k = specific rate constant. t = time elapsed since beginning of reaction.
a = stoichiometric coefficient of A in balanced overall equation.

31. Temperature: The Arrhenius Equation
Svante Arrhenius developed this relationship among (1) the temperature (T), (2) the activation energy (Ea), and (3) the specific rate constant (k).
If the Arrhenius equation is written for two temperatures, T2 and T1 with T2 >T1.

32. Temperature: The Arrhenius Equation
Subtract one equation from the other.
Rearrange and solve for ln k2/k1.

33. Temperature: The Arrhenius Equation
Consider the rate of a reaction for which Ea=50 kJ/mol, at 20oC (293 K) and at 30oC (303 K).
How much do the two rates differ?

34. Temperature: The Arrhenius Equation
For reactions that have an Ea»50 kJ/mol, the rate approximately doubles for a 100C rise in temperature, near room temperature.
Consider:
2 ICl(g) + H2(g) ® I2(g) + 2 HCl(g)
The rate-law expression is known to be R=k[ICl][H2].

35. Catalysts Homogeneous catalysts exist in same phase as the reactants.
Catalysts change reaction rates by providing an alternative reaction pathway with a different activation energy.
Homogeneous catalysts exist in same phase as the reactants.
Heterogeneous catalysts exist in different phases than the reactants.
Catalysts are often solids.

36. General rate expression
Differential rate law (with respect to concentration)
Rate = k [A]m[B]n[C]p
If zero order Rate = k[A]0 = k
first order Rate = k[A]1 = k[A]
second order Rate = k[A]2
Integrated rate law (with respect to time)
If zero order [A]0 - [A] = ak t
first order
second order
Kinetics
Initial rate
Instantaneous rate
Average rate
Rate of reaction
Order of reaction
Overall order of reaction
Catalyst
Four factors that affect the
rate of reaction
nature of reactant
concentration
temperature
presence of a catalyst
[A] vs t
ln [A] vs t
1/[A] vs t
[A]
ln [A]
1/[A] t t t

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45. 55. A reaction has the following experimental rate equation: Rate = k[A]2[B]. If the
concentration of A is doubled and the concentration of B is halved what happens
to the rate?
The following statements relate to the reaction with the following rate law:
Rate = k[H2][I2] H2(g) + I2(g) → 2 HI(g)
Determine which of the following statements are true. If a statement is false, indicate why it is incorrect.
The reaction must occur in a single step
This is a second-order reaction overall
Raising the temperature will cause the value of k to decrease
Raising the temperature lowers the activation energy for this reaction
If the concentration of both reactants are doubled, the rate will double
Adding a catalyst in the reaction will cause the initial rate to decrease

46. Describe each of the following statements as true or false
Describe each of the following statements as true or false. If false, rewrite the sentence to make it correct.
The rate determining elementary step in a reaction is the slowest step in a mechanism
It is possible to change the rate constant by changing the temperature
As a reaction proceeds at constant temperature, the rate remains constant
A reaction that is third order overall must involve more than one step
Reactions are faster at a higher temperature because activation energies are lower
Rate increase with increasing concentration of reactants because there are more collisions between reactant molecules
At higher temperatures a larger fraction of molecules have enough energy to get over the activation energy barrier
Catalyzed and uncatalyzed reactions have identical mechanisms