1. Chemical Kinetics
The area of chemistry concerned with the speeds, or rates, at which a chemical reaction occurs.
Chemical Kinetics

2. Reaction Rate
The reaction rate is the change in the concentration of a reactant or a product with time, (M/s or M . s-1), where M is molarity and s represents seconds.
Another way to represent rate is mol . L-1 s-1
Chemical Kinetics

3. Factors that Influence Reaction Rate
Under a given set of conditions, each reaction has its own characteristic rate, which is ultimately determined by the chemical nature of the reactants. (You will remember this from Chem I - potassium and water have a different rate of reaction than iron and oxygen.)
For a given reaction (using the same reactants), we can control four factors that affect its rate: the concentration of reactants, their physical state, the temperature at which the reaction occurs, and the use of a catalyst.
Chemical Kinetics

4. Concentration Rate collision frequency concentration
Molecules must collide in order to react. The more frequently they collide, the more often a reaction occurs. Thus, reaction rate is proportional to the concentration of reactant
Means “proportional to”
Rate collision frequency concentration
Therefore, if we increase the concentration…
we increase the collision frequency, which…
increases the rate
Chemical Kinetics

5. Physical State
Molecules must mix in order to collide. When reactants are in the same phase, as in aqueous solution, occasional stirring keeps them in contact. When they are in different phases, more vigorous mixing is needed. The more finely divided a solid or liquid reactant, the greater the surface are per unit volume, the more contact it makes with the other reactant, and the faster the reaction.
Chemical Kinetics

6. Temperature
Molecules must collide in order to react. Since the speed of a molecule depends on its temperature, more collisions will occur if the temperature is increased.
Speed of a molecule
Number of collisions
Chemical Kinetics

7. Temperature
Molecules must also collide with enough energy to react. Increasing the temperature increases the kinetic energy of the molecules, which in turn increases the energy of the collisions.
Chemical Kinetics

8. Temperature
Therefore, at a higher temperature, more collisions occur with enough energy to react. Thus, raising the temperature increases the reaction rate by increasing the number and especially the energy of the collisions.
Two familiar kitchen appliances employ this effect: a refrigerator slows down chemical processes that spoil food, whereas an oven speeds up other chemical processes to cook it.
Chemical Kinetics

9. Expressing Reaction Rate
Before we can deal quantitatively with the effects of concentration and temperature on reaction rate, we must be able to express the rate mathematically. A rate is a change in some variable per unit of time.
For example, the rate of motion of a car is the change of position of the car divided by time. A car that travels 57 miles in 60. minutes is traveling at…
57 miles/60. minutes = .95 miles/min
In the case of chemical reactions, the positions of the substances do not change over time, but their concentrations do.
Chemical Kinetics

10. We know that any reaction can be represented by the general equation
reactants  products
This equation tells us that during the course of a reaction, reactants are consumed while products are formed. As a result, we can follow the progress of a reaction by monitoring either the decrease in concentration of the reactants or the increase in concentration of the products.
Chemical Kinetics

11. The following figure shows the progress of a simple reaction in which A molecules are converted to B molecules:
A  B B A
Chemical Kinetics

12. The decrease in number of A molecules and the increase in the number of B molecules with time are shown below.
Chemical Kinetics

13. In general, it is more convenient to express the reaction rate in terms of the change in concentration with time. Thus, for the reaction A  B we can express the rate as:
D[A] Dt
[A]final – [A]initial
Rate = -
Because the concentration of A decreases during the time interval, D[A] is a negative quantity. The rate of a reaction is a positive quantity, so a minus sign is needed in the rate expression to make the rate positive.
Chemical Kinetics

14. or
D[B] Dt
Rate =
The rate of product formation does not require a minus sign because D[B] ([B]final – [B]initial) is a positive quantity (the concentration increases with time).
Chemical Kinetics

15. or D[B] Dt Rate = Rate = - D[A] Dt
where D[A] and D[B] are the changes in concentration (molarity) over a time period Dt.
These rates are average rates because they are averaged over a certain time period (Dt).
Chemical Kinetics

16. Reaction Rates and Stoichiometry
We have seen that for stoichiometrically simple reactions of the type A  B, the rate can either be expressed in terms of the decrease in reactant concentration with time, -D[A]/Dt, or the increase in product concentration with time, D[B]/Dt.
For more complex reactions, we must be careful in writing the rate expressions.
Consider the reaction
2A  B
Two moles of A disappear for each mole of B that forms
Chemical Kinetics

17. Another way to think of this is to say that the rate of disappearance of A is twice as fast as the rate of appearance of B. We write the rate as either
1 D[A]
2 Dt
D[B] Dt
Rate = - or
Rate =
For the reaction
2A  B
Chemical Kinetics

18. In general, for the reaction aA + bB  cC + dD The rate is given by
1 D[A]
a Dt
1 D[A]
a Dt
1 D[B]
b Dt
1 D[C]
c Dt
1 D[C]
c Dt
1 D[D]
d Dt
= - = =
Chemical Kinetics

19. Write the expression for the following reactions in terms of the disappearance of the reactants and the appearance of the products:
3O2(g)  2O3(g)
Rate = =
Chemical Kinetics

20. By definition, we know that to determine the rate of a reaction we have to monitor the concentration of the reactant (or product) as a function of time.
For reactions in solution, the concentration of a species can often be measured by spectroscopic means.
If ions are involved, the change in concentration can also be detected by an electrical conductance measurement.
Reactions involving gases are most conveniently followed by pressure measurements.
Chemical Kinetics

21. Reaction of Molecular Bromine and Formic Acid
In aqueous solutions, molecular bromine reacts with formic acid (HCOOH) as follows:
Br2(aq) + HCOOH(aq)  2Br-(aq) + 2H+(aq) + CO2(g)
Reddish-brown
colorless
The rate of Br2 disappearance can be determined by monitoring the color over time.
As the reaction proceeds, the color of the solution …
goes from brown to colorless
Chemical Kinetics

22. Reaction of Molecular Bromine and Formic Acid
Br2(aq) + HCOOH(aq)
2Br-(aq) + 2H+(aq) + CO2(g)
As the reaction proceeded, the concentration of Br2 steadily decreased and the color of the solution faded.
Chemical Kinetics

23. D[Br2] Dt Average rate = [Br2]t – [Br2]0 tfinal – tinitial
Measuring the change (decrease) in bromine concentration at some initial time ([Br2]0) and then at some other time, ([Br2]t) allows us to determine the average rate of the reaction during that interval:
D[Br2] Dt
Average rate =
[Br2]t – [Br2]0
tfinal – tinitial
Average rate =
Chemical Kinetics

24. Use the data in the following table to calculate the average rate over the first 50 second time interval.
Time
(s)
[Br2]
(M)
Rate
(M/s)
0.0
0.0120
4.20 x 10-5
50.0
0.0101
3.52 x 10-5
100.0
2.96 x 10-5
150.0
2.49 x 10-5
200.0
2.09 x 10-5
250.0
1.75 x 10-5
300.0
1.48 x 10-5
350.0
1.23 x 10-5
400.0
1.04 x 10-5
Print this chart
Chemical Kinetics

25. Now use the data in the same table to calculate the average rate over the first 100 second time interval.
Time
(s)
[Br2]
(M)
Rate
(M/s)
0.0
0.0120
4.20 x 10-5
50.0
0.0101
3.52 x 10-5
100.0
2.96 x 10-5
150.0
2.49 x 10-5
200.0
2.09 x 10-5
250.0
1.75 x 10-5
300.0
1.48 x 10-5
350.0
1.23 x 10-5
400.0
1.04 x 10-5
Chemical Kinetics

26. These calculations demonstrate that the average rate of the reaction depends on the time interval we choose.
By calculating the average reaction rate over shorter and shorter intervals, we can obtain the rate for a specific instant in time, which gives us the instantaneous rate of the reaction at that time.
Chemical Kinetics

27. The figure below shows the plot of [Br2] versus time, based on the data table given previously. Graphically, the instantaneous rate at 100 seconds after the start of the reaction is the slope of the line tangent to the curve at that instant.
Unless otherwise stated, we will refer to the instantaneous rate as simply “the rate”.
The instantaneous rate at any other time can be determined in a similar manner.
Chemical Kinetics

28. rate [Br2] rate = k[Br2] k, the Rate Constant
At a specific temperature, a rate constant (k) is a constant of proportionality between the reaction rate and the concentrations of reactants.
rate
[Br2]
Means “proportional to”
rate = k[Br2]
k is specific for a given reaction at a given temperature; it does not change as the reaction proceeds.
Chemical Kinetics

29. Rearrange the equation rate = k[Br2] To solve for k
Since reaction rate has the units M/s, and [Br2] is in M, the unit of k for this first order reaction is 1/s or s-1.
Rate
[Br2]
k =
Chemical Kinetics

30. Calculate the rate constant for the following reaction
Br2(aq) + HCOOH(aq)  2Br-(aq) + 2H+(aq) + CO2(g)
Time
(s)
[Br2]
(M)
Rate
(M/s)
0.0
0.0120
4.20 x 10-5
50.0
0.0101
3.52 x 10-5
100.0
2.96 x 10-5
150.0
2.49 x 10-5
200.0
2.09 x 10-5
250.0
1.75 x 10-5
300.0
1.48 x 10-5
350.0
1.23 x 10-5
400.0
1.04 x 10-5
k = rate/[Br2]
(s-1)
Chemical Kinetics

31. Because k is a constant (for this reaction at this specific temperature), it doesn’t matter which row we consider, so let’s consider the data at time 0.0 seconds…
Time
(s)
[Br2]
(M)
Rate
(M/s)
0.0
0.0120
4.20 x 10-5
50.0
0.0101
3.52 x 10-5
100.0
2.96 x 10-5
150.0
2.49 x 10-5
200.0
2.09 x 10-5
250.0
1.75 x 10-5
300.0
1.48 x 10-5
350.0
1.23 x 10-5
400.0
1.04 x 10-5
k = rate/[Br2]
(s-1)

32. To prove that k is a constant, calculate k at time 200.0 seconds
[Br2]
(M)
Rate
(M/s)
0.0
0.0120
4.20 x 10-5
50.0
0.0101
3.52 x 10-5
100.0
2.96 x 10-5
150.0
2.49 x 10-5
200.0
2.09 x 10-5
250.0
1.75 x 10-5
300.0
1.48 x 10-5
350.0
1.23 x 10-5
400.0
1.04 x 10-5
k = rate/[Br2]
(s-1)
The slight variations in the values of k are due to experimental deviations in rate measurements.
Chemical Kinetics

33. Filling in the rest of the table…
Time
(s)
[Br2]
(M)
Rate
(M/s)
0.0
0.0120
4.20 x 10-5
50.0
0.0101
3.52 x 10-5
100.0
2.96 x 10-5
150.0
2.49 x 10-5
200.0
2.09 x 10-5
250.0
1.75 x 10-5
300.0
1.48 x 10-5
350.0
1.23 x 10-5
400.0
1.04 x 10-5
k = rate/[Br2]
(s-1)
3.50 x 10-3
3.49 x 10-3
3.50 x 10-3
3.51 x 10-3
3.51 x 10-3
3.50 x 10-3
3.52 x 10-3
3.48 x 10-3
3.51 x 10-3
Chemical Kinetics

34. Units of the Rate Constant k for Several Overall Reaction Orders
Units of k (when t is seconds)
mol/L . s (or mol L-1 s-1) 1
1/s (or s-1) 2
L/mol . s (or L mol-1 s-1) 3
L2/mol2 . s (or L2 mol-2 s-1)
Chemical Kinetics

35. It is important to understand that k is NOT affected by the concentration of Br2.
The rate is greater at a higher concentration and smaller at a lower concentration of Br2, but the ratio of rate/[Br2] remains the same provided the temperature doesn’t change.
Chemical Kinetics

36. The Rate Law
The rate law expresses the relationship of the rate of a reaction to the rate constant (k) and the concentrations of the reactants raised to a power. For the general reaction
aA + bB  cC + dD
The rate law takes the form
Rate = k[A]x[B]y
Where x and y are numbers that must be determined experimentally.
Note – in general, x and y are NOT equal to the stoichiometric coefficients a and b from the overall balanced chemical equation. When we know the values of x, y and k, we can use the rate equation shown above to calculate the rate of the reaction, given the concentrations of A and B.
Chemical Kinetics

37. Rate = k[A]x[B]y
The reaction orders define how the rate is affected by the concentration of each reactant.
This reaction is xth order in A, yth order in B.
Chemical Kinetics

38. The following rate law was determined for the formation of nitrogen trioxide and molecular oxygen from nitrogen dioxide and ozone
Rate = k[NO2][O3]
How would the rate of this reaction be affected if the concentration of NO2 increased from 1.0 M to 2.0 M?
This reaction is first order with respect to both NO2 and O3. This means that doubling the concentration of either reactant would double the rate of the reaction.
(2)1 = 2
How many times greater the concentration is
What the order is for that reactant
How many times greater the rate of the reaction will be

39. The reaction between nitrogen monoxide and molecular oxygen is described by a different rate law.
Rate = k [NO]2[O2]
How would the rate of this reaction be affected if the concentration of NO increased from 1.0 M to 2.0 M?
Chemical Kinetics

40. Rate = k [NO]2[O2]
How would the rate of this same reaction be affected if the concentration of NO increased from 1.0 M to 5.0 M?
Chemical Kinetics

41. The overall reaction order is x + y.
The exponents x and y specify the relationships between the concentrations of reactants A and B and the reaction rate. Added together, they give us the overall reaction order, defined as the sum of the powers to which all reactant concentrations appearing in the rate law are raised. For the equation
Rate = k[A]x[B]y
The overall reaction order is x + y.
Chemical Kinetics

42. (CH3)3CBr(l) + H2O(l)  (CH3)3COH(l) + HBr(aq)
For the following reaction
(CH3)3CBr(l) + H2O(l)  (CH3)3COH(l) + HBr(aq)
The rate law has been found to be
rate = k[(CH3)3CBr]
This reaction is first order in 2-bromo-2-methylpropane.
Note that the concentration of H2O does not even appear in the rate law. Thus, the reaction is zero order with respect to H2O. This means that the rate does not depend on the concentration of H2O; we could also write the rate law for this reaction as
rate = k[(CH3)3CBr][H2O]0
What is the overall order of this reaction?
1st order overall (1+0=1)
Chemical Kinetics

43. CHCl3(g) + Cl2(g)  CCl4(g) + HCl(g)
Reaction orders are usually positive integers or zero, but they can also be fractional or negative. In the reaction
CHCl3(g) + Cl2(g)  CCl4(g) + HCl(g)
A fractional order appears in the rate law:
rate = k[CHCl3][Cl2]1/2
This order means that the reaction depends on the square root of the Cl2 concentration. If the initial Cl2 concentration is increased by a factor of 4, for example, the rate increases by V4 (= 2), therefore the rate would double.
Chemical Kinetics

44. A negative exponent means that the reaction rate decreases when the concentration of that component increases. Negative orders are often seen for reactions whose rate laws include products. For example, in the atmospheric reaction
2O3(g) O2(g)
The rate law has been shown to be
Rate = k[O3]2[O2]-1 ; or [O3]2
rate = k
[O2]
If the [O2] doubles, the reaction proceeds half as fast.
Chemical Kinetics

45. To see how to determine the rate law of a reaction, let us consider the reaction between fluorine and chlorine dioxide:
F2(g) + 2ClO2(g)  2FClO2(g)
Chemical Kinetics

46. One way to study the effect of reactant concentration on reaction rate is to determine how the initial rate depends on the starting concentrations. It is preferable to measure the initial rates because as the reaction proceeds, the concentrations of the reactants decrease and it may become difficult to measure the changes accurately. Also, as the reaction continues, the product concentrations increase,
products  reactants
so the reverse reaction becomes increasingly likely. Both of these complications are virtually absent during the earliest stages of the reaction.
Chemical Kinetics

47. The following table shows three rate measurements for the formation of FClO2.
Initial Rate
(M/s)
0.10
0.010
1.2 x 10-3
0.040
4.8 x 10-3
0.20
2.4 x 10-3
Chemical Kinetics

48. Looking at entries 1 and 3, we see that as we double [F2]0 while holding [ClO2]0 constant, the reaction rate doubles. Thus the rate is directly proportional to [F2], and the reaction is first order with respect to F2.
[F2]0
(M)
[ClO2]0
Initial Rate
(M/s)
0.10
0.010
1.2 x 10-3
0.040
4.8 x 10-3
0.20
2.4 x 10-3
Chemical Kinetics

49. Similarly, the data in entries 1 and 2 show that as we quadruple [ClO2] while holding [F2] constant, the rate increases by four times, so the rate is also directly proportional to [ClO2], making the reaction 1st order with respect to [ClO2]
[F2]
(M)
[ClO2]
Initial Rate
(M/s)
0.10
0.010
1.2 x 10-3
0.040
4.8 x 10-3
0.20
2.4 x 10-3
Chemical Kinetics

50. We can summarize our observations by writing the rate law as
Rate = k[F2][ClO2]
Because both [F2] and [ClO2] are raised to the first power, the reaction is second order overall.
Chemical Kinetics

51. What is the rate constant for this reaction at this temperature?
Chemical Kinetics

52. 2NO(g) + 2H2(g)  N2(g) + 2H2O(g)
The reaction of nitric oxide with hydrogen at 1280 oC is
2NO(g) + 2H2(g)  N2(g) + 2H2O(g)
From the following data that was collected experimentally at this temperature, determine the rate law and calculate the rate constant.
Experiment
[NO]
[H2]
Initial Rate (M/s) 1
5.0 x 10-3
2.0 x 10-3
1.3 x 10-5 2
10.0 x 10-3
5.0 x 10-5 3
4.0 x 10-3
10.0 x 10-5
Chemical Kinetics

53. rate law… Experiment [NO] [H2] Initial Rate (M/s) 1 5.0 x 10-3
Chemical Kinetics

54. rate constant… Experiment [NO] [H2] Initial Rate (M/s) 1 5.0 x 10-3
Chemical Kinetics

55. The following points summarize our discussion of the rate law:
Rate laws are ALWAYS determined experimentally.
Reaction order should be defined in terms of reactant (not product) concentrations.
The order of a reactant is not related to the stoichiometric coefficient of the reactant in the overall balanced equation.
Chemical Kinetics

56. Relationship Between Reactant Concentration and Time
Rate law expressions enable us to calculate the rate of a reaction from the rate constant and reactant concentrations.
The rate laws can also be used to determine the concentrations of reactants any time during the course of a reaction.
Chemical Kinetics

57. First-order reactions
A first order reaction is a reaction whose rate depends on the reactant concentration raised to the first power.
rate = k[A]
In a first-order reaction of the type
A  product
Rate can be expressed as
rate = -
D[A] Dt
Chemical Kinetics

58. We can write the integrated rate law for a first order reaction as
Through the methods of calculus, this expression is integrated over time to obtain the integrated rate law for a first-order reaction:
[A]0
[A]t ln
= kt ln
[A]0
[A]t
Since
= ln[A]0 – ln[A]t
We can write the integrated rate law for a first order reaction as
ln[A]0 – ln[A]t = kt
Chemical Kinetics

59. Cyclobutane decomposes at 1000 oC to two moles of ethylene (C2H4) with a very high rate constant, 87 s-1.
If the initial concentration of cyclobutane is 2.00 M, what is the concentration after s?
87 s-1
Chemical Kinetics

60. What percent of the cyclobutane has decomposed in this time?
Chemical Kinetics

61. Reaction half-life (t1/2)
The t1/2 of a reaction is the time required to reach half the initial reactant concentration.
For a first order reaction, the formula for determining t1/2 is
Note – t1/2 of a first order reaction is constant, it is independent of reactant concentration!
ln 2 k
0.693 k
t1/2 = =
Chemical Kinetics

62. A plot of [N2O5] vs time for three half-lives.
0.060
0.050
0.040
[N2O5]
0.030
0.020
0.010
0.000 24 48 72
Time (min)
Chemical Kinetics

63. Determining t1/2 for a first-order reaction
Cyclopropane is the smallest cyclic hydrocarbon. Because its 60o bond angles allow only poor orbital overlap, its bonds are weak. As a result, it is thermally unstable and rearranges to propene at 100 oC via the following reaction:
CH2
 CH3 – CH = CH2
H2C
CH2
Chemical Kinetics

64. The rate constant is 9.2 s-1. How long does it take for the initial concentration of cyclopropane to decrease by one-half?
Chemical Kinetics

65. Second-order reactions
For a general second-order rate equation, the expression including time can become quite complex, so let’s consider only the simplest case, one in which the rate law contains only one reactant
2A  product
The integrated rate law for a second-order reaction involving one reactant: 1
[A]t
[A]0 -
= kt
Chemical Kinetics

66. rate = k[HI]2 and k = 2.4 x 10-21 L/mol . s
At 25 oC, hydrogen iodide breaks down very slowly to hydrogen and iodine according to the following:
rate = k[HI]2 and k = 2.4 x L/mol . s
If mol HI(g) is placed in a 1.0 L container, how long will it take for the concentration of HI to reach mol/L?
Chemical Kinetics

67. In contrast to the half-life of a first-order reaction, the half-life of a second-order reaction DOES depend on reactant concentration: 1
k[A]0
t1/2 =
Chemical Kinetics

68. rate = k[AB]2 and k = 0.20 L/mol . s
In the simple decomposition reaction
AB(g)  A(g) + B(g)
rate = k[AB]2 and k = 0.20 L/mol . s
How long will it take for [AB] to reach half of its initial concentration of 1.50 M?
Chemical Kinetics

69. Determining the Reaction Order from the Integrated Rate Law
Suppose you don’t know the rate law for a reaction and don’t have the initial rate data needed to determine the reaction orders. Another method for finding reaction orders is a graphical technique that uses concentration-time data directly.
Chemical Kinetics

70. To find the reaction order from the concentration-time data, some trial-and-error graphical plotting is required:
If you obtain a straight line when you plot [reactant] vs. time, the reaction is zero order with respect to that reactant.
If you obtain a straight line when you plot ln[reactant] vs. time, the reaction is first order with respect to that reactant.
If you obtain a straight line when you plot 1/[reatant] vs. time, the reaction is second order with respect to that reactant.
Chemical Kinetics

71. This graph shows that the rate is ___ order with respect to A 2nd
Chemical Kinetics

72. Graphical determination of the reaction order for the decomposition of N2O5.
A plot of 1/[N2O5] vs t is curved, indicating that the reaction IS NOT second order in N2O5.
Time
[N2O5]
ln[N2O5]
1/[N2O5]
0.0165
-4.104
60.6 10
0.0124
-4.390
80.6 20
0.0093
-4.68
1.1 x 102 30
0.0071
-4.95
1.4 x 102 40
0.0053
-5.24
1.9 x 102 50
0.0039
-5.55
2.6 x 102 60
0.0029
-5.84
3.4 x 102
Chemical Kinetics

73. Graphical determination of the reaction order for the decomposition of N2O5.
A plot of ln[N2O5] vs t gives a straight line, indicating the reaction IS first order in N2O5
Time
[N2O5]
ln[N2O5]
1/[N2O5]
0.0165
-4.104
60.6 10
0.0124
-4.390
80.6 20
0.0093
-4.68
1.1 x 102 30
0.0071
-4.95
1.4 x 102 40
0.0053
-5.24
1.9 x 102 50
0.0039
-5.55
2.6 x 102 60
0.0029
-5.84
3.4 x 102
Chemical Kinetics

74. The Collision Theory of Chemical Kinetics
The kinetic molecular theory of gases postulates that gas molecules frequently collide with one another. Therefore, it seems logical to assume – and it is generally true - that chemical reactions occur as a result of collisions between reacting molecules.
In terms of the collision theory of chemical kinetics, then, we expect the rate of a reaction to be directly proportional to the frequency of the collisions (number of molecular collisions per second).
Rate =
Number of collisions s
Chemical Kinetics

75. The Collision Theory of Chemical Kinetics
This simple relationship explains the dependence of reaction rate on concentration.
Increasing the concentration increases the likelihood that molecules will collide.
Rate =
Number of collisions s
Chemical Kinetics

76. However, not all collisions lead to reactions.
Energetically speaking, there is some minimum collision energy below which no reaction occurs. Any molecule in motion possess kinetic energy; the faster it moves, the greater its kinetic energy. When molecules collide, part of their kinetic energy is converted to vibrational energy. If the initial kinetic energies are large enough, the colliding molecules will vibrate so strongly that some of the chemical bonds will break. This bond fracture is the first step toward product formation. If the initial kinetic energies are too small, the molecules will merely bounce off each other intact, and no change results from the collision.
Chemical Kinetics

77. We postulate that in order to react, the colliding molecules must have a total kinetic energy equal to or greater than the activation energy, (Ea), which is defined as the minimum amount of energy required to initiate a chemical reaction.
Chemical Kinetics

78. Molecules must also be oriented in a favorable position – one that allows the bonds to break and atoms to rearrange.
Chemical Kinetics

79. NO(g) + NO3(g)  2NO2(g)
The picture to the right shows a few of the possible collision orientations for this simple gaseous reaction:
Of the five collisions shown, only one has an orientation in which the N of NO collides with an O of NO3. Actually, the probability factor (p) for this reaction is 0.006; only 6 collisions in every thousand have an orientation that leads to a reaction.
Chemical Kinetics

80. Collisions between individual atoms have p values near 1: almost no matter how they hit, they react. In such cases, the rate constant depends only on the frequency and energy of the collisions.
At the other extreme are biochemical reactions, in which the reactants are often two small molecules that can react only when they collide with a specific tiny region of a giant molecule-a protein or nucleic acid. The orientation factor for such reactions is often less than 10-6: fewer than one in a million sufficiently energetic collisions leads to product formation.
Chemical Kinetics

81. When reactants collide at the proper angle with energy equal to the activation energy (Ea), they undergo an extremely brief interval of bond disruption and bond formation called a transition state.
During this transition state, the reactants form a short-lived complex that is neither reactant nor product, but has partial bonding characteristics of both. This transitional structure is called an activated complex.
Endothermic/Exothermic (choose one)
Chemical Kinetics

82. An activated complex is a highly unstable species with a high potential energy. (It was energized by the particle collision.) Once formed, it will break up almost immediately.
The activated complex exists along the reaction pathway at the point where the energy is greatest – at the peak indicated by the activation energy.
Activation energy is the energy required to achieve the transition state and form the activated complex
Chemical Kinetics

83. We can think of activation energy, (Ea), as a barrier that prevents less energetic molecules from reacting…
…because only the molecules who have enough kinetic energy to exceed the activation energy can take part in the reaction.
Chemical Kinetics

84. Maxwell Boltzmann Diagram
Because the number of reactant molecules in an ordinary reaction is very large, the speeds, and hence also the kinetic energies of the molecules, vary greatly.
Chemical Kinetics

85. Normally, only a small fraction of the colliding molecules -- the fastest-moving ones – have enough kinetic energy to exceed the activation energy.
At higher temperature, more molecules can surpass the activation energy, therefore, the rate of product formation is greater at the higher temperature.
Chemical Kinetics

86. With very few exceptions, reaction rates increase with increasing temperature.
As a general rule of thumb, you can expect a 10 oC increase in temperature to result in a doubling of the reaction rate.
Chemical Kinetics

87. Arrhenius Equation k = Ae-Ea/RT
The dependence of the rate constant of a reaction on temperature can be expressed by the following equation, known as the Arrhenius equation:
k = Ae-Ea/RT
Where Ea is the activation energy (in J/mol), R the gas constant (8.314 J/K . mol), T the absolute temperature, and e the base of the natural logarithm scale. The quantity A represents the collision frequency and is called the frequency factor.
Chemical Kinetics

88. Frequency Factor (A)
The frequency factor is the product of the collision frequency (Z) and an orientation factor (p) which is specific for each reaction. The factor p is related to the structural complexity of the colliding particles. You can think of it as the ratio of effectively oriented collisions to all possible collisions.
In the activation energy problems we are solving, the actual value of A need not be known because A can be treated as a constant for a given reacting system over a fairly wide temperature range.
Chemical Kinetics

89. As the activation energy increases, k decreases,
k = A e-Ea/RT
As the activation energy increases, k
decreases,
and as k decreases, rate
decreases
As the temperature increases, k
increases
And as k increases, rate
increases
Chemical Kinetics

90. You can derive the following equation from the Arrhenius equation:
R T
ln k =
+ ln A
Chemical Kinetics

91. slope (m) is equal to –Ea/R -Ea 1 R T ln k = + ln Amx
+ b
Y =
Thus, a plot of ln k versus 1/T gives a straight line whose
slope (m) is equal to –Ea/R and whose
intercept (b) with the y-axis is ln A.
Chemical Kinetics

92. See the following example…
Given k and temperature, you can use your graphing calculator to determine the activation energy of a reaction.
See the following example…
Chemical Kinetics

93. CH3CHO(g)  CH4(g) + CO(g)
The rate constants for the decomposition of acetaldehyde
CH3CHO(g)  CH4(g) + CO(g)
were measured at five different temperatures. The data are shown in the following table.
Chemical Kinetics

94. Enter these values into L1 and L2 of your graphing calculator.
T (K)
700
730
760
790
810 k
0.011
0.035
0.105
0.343
0.789
Chemical Kinetics

95. 1/T
1.43 x 10-3
1.37 x 10-3
1.32 x 10-3
1.27 x 10-3
1.23 x 10-3
ln k
-4.51
-3.35
-2.254
-1.070
0.237
To determine the activation energy, we need to graph 1/T on the x axis (L3) and ln k on the y-axis (L4).
The slope of the line can be determined using linear regression.
Chemical Kinetics

96. Chemical Kinetics

97. An equation relating the rate constants k1 and k2 at temperature T1 and T2 can be used to calculate the activation energy or to find the rate constant at another temperature if the activation energy is known. k2 k1 Ea R 1 ln =
T2 T1
Chemical Kinetics

98. The rate constant of a first-order reaction is 3
The rate constant of a first-order reaction is 3.46 x 10-2 s-1 at 298 K. What is the rate constant at 350 K if the activation energy for the reaction is 50.2 kJ/mol?
Chemical Kinetics

99. Reaction Mechanisms
As we mentioned earlier, an overall balanced chemical equation does not tell us much about how a reaction actually takes place. In many cases, it merely represents the sum of several elementary steps, or elementary reactions, that represent the progress of the overall reaction at the molecular level. The term for the sequence of elementary steps that leads to product formation is reaction mechanism.
The reaction mechanism is comparable to the route of travel followed during a trip; the overall chemical equation specifies only the origin and the destination.
Chemical Kinetics

100. As an example of a reaction mechanism, let us consider the reaction between nitrogen monoxide and oxygen:
2NO(g) + O2(g)  2NO2(g)
We know that the products are not formed directly from the collision of two NO molecules with an O2 molecule because N2O2 is detected during the course of the reaction.
Chemical Kinetics

101. Let us assume that the reaction actually takes place via two elementary steps as follows:
NO(g) + NO(g)  N2O2(g)
N2O2(g) + O2(g)  2NO2(g)
Elementary Step 2
Overall reaction
2NO + N2O2 + O2  N2O NO2
Each of the elementary steps listed above is called a bimolecular reaction because each step involves two reactant molecules. A step that just involves one reactant molecule is a unimolecular reaction. Very few termolecular reactions, reactions that involve the participation of three reactant molecules in one elementary step, are known, because the simultaneous encounter of three molecules is a far less likely event than a bimolecular collision. (There are no known examples of reactions involving the simultaneous encounter of four molecules.)
Chemical Kinetics

102. Elementary Step 1 NO(g) + NO(g)  N2O2(g) N2O2(g) + O2(g)  2NO2(g)
Overall reaction
2NO + N2O2 + O2  N2O NO2
Species such as N2O2 are called intermediates because they appear in the mechanism of the reaction (that is, in the elementary steps) but not in the overall balanced equation. Keep in mind that an intermediate is always formed in an early elementary step and consumed in a later elementary step. Note – an intermediate ≠ activated complex!
Chemical Kinetics

103. You can propose a mechanism for a reaction if you consider…
The elementary steps in a multistep reaction mechanism must always add to give the balanced chemical equation of the overall process. (Any intermediates that are formed in earlier steps must be consumed in later steps.)
Unimolecular and bimolecular reactions are more common than termolecular reactions.
The rate of the overall reaction is limited by the rate of the slowest elementary step, (For that reason, the slowest elementary step is typically called the rate-determining step.
Chemical Kinetics

104. Step 1: NO(g) + F2(g) NOF2(g) Step 2: NOF2(g) + NO(g)  2NOF(g)
Gaseous nitrogen monoxide reacts with fluorine gas to produce nitrogen hypofluorite (NOF). The intermediate product NOF2(g) has been isolated as an intermediate in this reaction. Propose a two step mechanism consistent with this intermediate product.
Step 1: NO(g) + F2(g) NOF2(g)
Step 2: NOF2(g) + NO(g)  2NOF(g)
The elementary steps add to give the overall balanced chemical equation
The first and the second step are bimolecular
Chemical Kinetics

105. Even when a proposed mechanism is consistent with the rate law, later experimentation may show it to be incorrect or only one of several alternatives.
Chemical Kinetics

106. Suppose we have the following elementary step: A  products
Knowing the elementary steps of a reaction enables us to propose a rate law.
Suppose we have the following elementary step:
A  products
Because there is only one reactant molecule present, this is a/n ___molecular reaction. It follows that the larger the number of A molecules present, the faster the rate of product formation.
Thus the rate of a unimolecular reaction is directly proportional to the concentration of A, or is first order in A:
Rate = k[A]
uni
Chemical Kinetics

107. For a bimolecular elementary reaction involving A and B molecules
A + B  product
the rate of product formation depends on how frequently A and B collide, which in turn depends on the concentration of A and B. Thus we can express the rate as
Rate = k[A][B]
Chemical Kinetics

108. Similarly, for a bimolecular elementary reaction of the type
A + A  products or
2A  products
the rate becomes
Rate = k[A]2
Chemical Kinetics

109. Rate Laws for Elementary Steps
Molecularity
Rate Law
A  product
2A  product
A + B  product
2A + B  product
unimolecular
rate = k[A]
bimolecular
rate = k[A]2
bimolecular
rate = k[A][B]
termolecular
rate = k[A]2[B]
Chemical Kinetics

110. Remember, when we study a reaction that has more than one elementary step, the rate law for the overall process is given by the rate-determining step, which is the slowest step in the sequence of steps leading to product formation.
Chemical Kinetics

111. Experimentally the rate law is found to be Rate = k[N2O]
Example: The gas-phase decomposition of dinitrogen monoxide (N2O) is believed to occur via two elementary steps:
Step 1: N2O N2 + O
Step 2: N2O + O  N2 + O2
Experimentally the rate law is found to be
Rate = k[N2O]
Write the equation for the overall reaction.
Identify the intermediates.
(c) What can you say about the relative rates of steps 1 and 2?
Chemical Kinetics

112. Example 2: Hydrogen Peroxide Decomposition
Does the decomposition of hydrogen peroxide occur in a single step?
The overall reaction is
2H2O2(aq)  2H2O(l) + O2(g)
By experiment, the rate law is found to be
Rate = k[H2O2][I-]
Chemical Kinetics

113. From this alone you can see that H2O2 decomposition does not occur in a single elementary step corresponding to the overall balanced equation. If it did, the rate law would be
Rate = [H2O2]2
or in other words, the reaction would be second order in H2O2. Remember, the experimentally determined rate law for this reaction was shown to be
Rate = k[H2O2][I-]
Chemical Kinetics

114. Catalysis 2H2O2(aq)  2H2O(l) + O2(g) Rate = k[H2O2][I-]
For the decomposition of hydrogen peroxide we see that the reaction rate depends on the concentration of iodide ions even though I- does not appear in the overall equation. I- is a catalyst for this reaction, a substance that increases the rate of a chemical reaction without itself being consumed.
Chemical Kinetics

115. A catalyst exists before the reaction occurs and can be recovered and reused after the reaction is complete. This is the opposite of intermediates, which are produced in one step of a mechanism and consumed in another.
Chemical Kinetics

116. In many cases, a catalyst increases the rate by providing a set of elementary steps with more favorable kinetics than those that exist in its absence.
In other words, a catalyst typically lowers the activation energy for the reaction by forming an activated complex that has less potential energy.
Chemical Kinetics

117. Forward reaction Reverse reaction
A + B
C + D
Forward reaction
A + B
C + D
Reverse reaction
Because the activation energy for the reverse reaction is also lowered, a catalyst enhances the rates of the forward and reverse reaction equally.
Chemical Kinetics

118. There are three general types of catalysis: heterogeneous catalysis, homogeneous catalysis, and enzyme catalysis.
Chemical Kinetics

119. In heterogeneous catalysis the reactants and the catalyst are in different phases. Usually the catalyst is a solid and the reactants are either gases or liquids. Heterogeneous catalysis is by far the most important type of catalysis in industrial chemistry, especially in the synthesis of many key chemicals.
Chemical Kinetics

120. N2(g) + 3H2(g)  2NH3(g) DH = -92.6 kJ
Ammonia is an extremely valuable inorganic substance used in the fertilizer industry and many other applications.
N2(g) + 3H2(g)  2NH3(g) DH = kJ
This reaction is extremely slow at room temperature, and although raising the temperature accelerates the above reaction, it also promotes the decomposition of NH3 molecules into N2 and H2, thus lowering the yield of NH3.
In 1905, after testing literally hundreds of compounds at various temperatures and pressures, Fritz Haber discovered that iron plus a few percent of oxides of potassium and aluminum catalyze the reaction. This procedure is known as the Haber process.
Chemical Kinetics

121. Haber Process
First the H2 and the N2 molecules bind to the surface of the catalyst. This interaction weakens the covalent bonds within the molecules and eventually causes the molecules to dissociate. The highly reactive H and N atoms combine to form NH3 molecules, which then leave the surface.
Chemical Kinetics

122. Nitric acid is one of the most important inorganic acids
Nitric acid is one of the most important inorganic acids. It is used in the production of fertilizers, dyes, drugs, and in many other products. The major industrial method of producing nitric acid is the Ostwald process. The starting materials, ammonia and molecular oxygen, are heated in the presence of a platinum-rhodium catalyst to about 800 oC.
Chemical Kinetics

123. In homogeneous catalysis, the reactants are dispersed in a single phase, usually liquid. Acid and base catalyses are the most important types of homogeneous catalysis in liquid solution.
Homogeneous catalysis can also take place in the gas phase. A well-known example of catalyzed gas-phase reactions is the lead chamber process, which for many years was the primary method of manufacturing sulfuric acid.
Chemical Kinetics

124. Homogeneous catalysis has several advantages over heterogeneous catalysis. For one thing, the reactions can often be carried out under atmospheric conditions, thus reducing production costs and minimizing the decomposition of products at high temperatures. In addition, homogeneous catalysts can be designed to function selectively for a particular type of reaction, and homogeneous catalysts cost less than the precious metals (for example, platinum and gold) used in heterogeneous catalysis.
Chemical Kinetics

125. Of all the intricate processes that have evolved in living systems, none is more striking or more essential than enzyme catalysis.
Enzymes are biological catalysts. Enzymes can increase the rate of a biochemical reaction by a factor ranging from 106 to 1012 times!
An enzyme acts only on certain molecules, called substrates while leaving the rest of the system unaffected. It has been estimated that an average living cell may contain some 3000 different enzymes, each of them catalyzing a specific reaction in which a substrate is converted into the appropriate products.
Chemical Kinetics

126. An enzyme is typically a large protein molecule that contains one or more active sites where interactions with substrates takes place. These sites are structurally compatible with specific substrate molecules, in much the same way as a key fits a particular lock. In fact, the notion of a rigid enzyme structure that binds only to molecules whose shape exactly matches that of the active site was the basis of an early theory of enzyme catalysis, the so-called lock-and-key theory.
Chemical Kinetics

127. This theory accounts for the specificity of enzymes, but it contradicts research evidence that a single enzyme binds to substrates of different sizes and shapes.
Chemists now know that an enzyme molecule (or at least its active site) has a fair amount of structural flexibility and can modify its shape to accommodate more than one type of substrate.
Chemical Kinetics

128. The End
Chemical Kinetics