1. Chapter 14
Chemical Kinetics

2. Reaction Rate
The rate of a chemical reaction is measured as the decrease in the concentration of a reactant or the increase in the concentration of a product in a unit of time.

3. Reaction Rate
The rate of a chemical reaction is measured as the decrease in the concentration of a reactant or the increase in the concentration of a product in a unit of time.
∆[ ]
Rate =
∆time

4. Reaction Rate
The rate of a chemical reaction is measured as the decrease in the concentration of a reactant or the increase in the concentration of a product in a unit of time.
∆[ ]
What units would we use for rate?
Rate =
∆time

5. Reaction Rate
The rate of a chemical reaction is measured as the decrease in the concentration of a reactant or the increase in the concentration of a product in a unit of time.
∆[ ]
Rate =
∆time
2H2O2(aq) → 2H2O(l) + O2(g)

6. Reaction Rate
The rate of a chemical reaction is measured as the decrease in the concentration of a reactant or the increase in the concentration of a product in a unit of time.
∆[ ]
Rate =
∆time
2H2O2(aq) → 2H2O(l) + O2(g)
How could the rate be expressed for this reaction in terms of H2O2?

7. 2H2O2(aq) → 2H2O(l) + O2(g)

8. 2H2O2(aq) → 2H2O(l) + O2(g)

9. What is the rate of the reaction from 0s to 2.16 x 104s?
2H2O2(aq) → 2H2O(l) + O2(g)
What is the rate of the reaction from 0s to 2.16 x 104s?

10. 1.16 x 10-5 mol O2 L-1 s-1 2H2O2(aq) → 2H2O(l) + O2(g)
What is the average rate of appearance of O2 from 0s to 2.16 x 104s?
1.16 x 10-5 mol O2 L-1 s-1

11. General Rate of Reaction
a A + b B → c C + d D
Rate of reaction = rate of disappearance of reactants or
Rate of reaction = rate of appearance (formation) of products
We can use the coefficients in the equation to compare the reaction rates for all the substances in the reaction.

12. 15-1 The Rate of a Chemical Reaction
Rate is change of concentration with time.
2 Fe3+(aq) + Sn2+(aq) → 2 Fe2+(aq) + Sn4+(aq)
t = 38.5 s [Fe2+] = M
∆t = 38.5 s ∆[Fe2+] = ( – 0) M
Rate of formation of Fe2+= = = 2.6 x 10-5 M s-1
Δ[Fe2+] Δt M
38.5 s

13. Rates of Chemical Reaction
2 Fe3+(aq) + Sn2+(aq) → 2 Fe2+(aq) + Sn4+(aq)
Rate of formation of Fe2+ = 2.6 x 10-5 mol L-1 s-1
What is the rate of formation of Sn4+?
1.3 x 10-5 mol Sn4+ L-1 s-1
What is the rate of disappearance of Fe3+?
2.6 x 10-5 mol Fe3+ L-1 s-1

15. What does the slope of the line represent?

16. What is the concentration at 100s for the reaction: 2H2O2(aq) → 2H2O(l) + O2(g)? Given:
Rate = 1.7 x 10-3 M s-1 Δt =
Δ[H2O2]
[H2O2]i = 2.32 M
Δ[H2O2] = (1.7 x 10-3 M s-1) (∆t)
∆[H2O2] = (1.7 x 10-3 M s-1)(100 s) = 0.17M
= 2.32 M M
[H2O2]100 s
= 2.15 M

17. What does it mean when the rate of a reaction reaches zero?
For a normal reaction it means that one or more of the reactants are used up and the reaction has stopped.
For a reversible reaction it means that the reaction has reached equilibrium.

18. Factors Affecting Reaction Rates
The nature of the reacting substances.

19. Factors Affecting Reaction Rates
2. The state of subdivision of the reacting substances (surface area).

20. Lycopodium Powder Explosion ≈ 30s

21. Factors Affecting Reaction Rates
3. The temperature of the reacting substances.

22. Factors Affecting Reaction Rates
4. The concentration of the reacting substances.
(Except in zero order reactions)
Air (21% oxygen) % oxygen

23. Factors Affecting Reaction Rates
The presence of a catalyst.
Catalysts speed up reactions but are left unchanged by the reaction.

24. Rate constant = k (is “k” really constant?)
The Rate Law
a A + b B …. → g G + h H ….
Rate = k [A]m[B]n ….
Rate constant = k (is “k” really constant?)
m and n are usually small whole numbers but may be fractional, negative or zero. They are often not related to a and b.
The larger k, the faster the reaction. k depends on temperature, concentration of catalyst and the specific reaction.
Order of A = m Order of B = n
Overall order of reaction = m + n + ….

25. Temperature and Rate
Generally, as temperature increases, so does the reaction rate.
This is because k is temperature dependent.
Therefore the temperature dependence of reaction rates is contained in the temperature dependence of the rate constant.

26. Temperature dependence of “k”. . . . .
The rate constant (k) is actually only constant if all you are changing is the concentration of the reactants in a reaction.
If you change the temperature, add a catalyst, or change the reaction the rate constant changes.

27. Concentration and Rate Summary
After finding the trials to compare:
A reactant is zero order if the change in concentration of that reactant produces no effect on the rate.
A reaction is first order if doubling the concentration of that reactant causes the rate to double.
A reactant is nth order if doubling the concentration of that reactant causes an 2n increase in rate.
Note that the rate constant does not depend on concentration.

28. Use the data provided to write the rate law and indicate the order of the reaction with respect to HgCl2 and C2O42- and also the overall order of the reaction.

29. First determine the order of HgCl2

30. Next determine the order of C2O42-

31. Now write the rate law and determine the order of the reaction.

32. Calculate the rate constant “k” and its units.
Initial rate of disappearance HgCl2
mol L-1 min-1

33. What is the average rate of disappearance of C2O42- in trial 1?
Initial rate of disappearance HgCl2
mol L-1 min-1

34. Use the data provided to write the rate law and indicate the order of the reaction with respect to NO2 and CO (support your answers). Also give the overall order of the reaction.

36. Calculate the rate constant “k” and its units.

37. What is the average rate of disappearance of CO in trial 2?

38. How do we make these charts?
Initial rate of disappearance HgCl2
mol L-1 min-1
Rates can be measured experimentally using a variety techniques:
moniter pH changes
Titrations
Change in volume or mass (gas production)
Basically we can use any method to follow a reaction that produces a measurable change.

39. How do we make these charts?
Initial rate of disappearance HgCl2
mol L-1 min-1
One important method involves the spectroscopic determination of concentration through Beer’s Law.

40. Using Beer’s Law to Determine [ ] vs. time.
For each trial, the reactants are mixed and the reaction mixture is transferred into a test tube or cuvette.
Without any delay, the reaction vessel is placed into a spectrophotometer. The absorbance data is then collected at the wavelength of maximum absorbance as a function of time.
This absorbance data is then converted to concentration data using Beer’s Law: A = ɛ l c

41. Fe(s)+CuSO4(aq)→Fe2SO4(aq)+Cu(s)
The solution gradually gets paler as the concentration of copper sulfate decreases and the concentration of iron sulfate increases.
Concentration of copper sulfate solution 1M
0.8M
0.6M
0.4M
0.2M 0s
30s
90s
200s
500s

42. Using Beer’s Law to Determine [ ] vs. time.
A graph of concentration vs. time can be prepared and then used to experimentally determine the rate.

43. What does this tangent allow us to measure?

44. Half Life of a First Order Reaction
Half-life is the time required to convert one half of a reactant to product.
For first-order reactions, half-life is often used as a representation for the rate constant.
This is because the half-life of a first-order reaction and the rate constant are inversely proportional, and the half-life is independent of concentration.
t½ = / k

45. Radioactivity
Radioactive decay is an example of a first order process.
Radioactive decay is the spontaneous breakdown of unstable atoms into more stable atoms with the simultaneous emission of particles and rays.
Radioactive decay occurs at a constant rate that is a first order process.

46. Radioactivity and Half - Life
The half-life of carbon-14 is 5730 years.
How old is a bone that has about 12.5% of the carbon-14 that a living organism would have in it?

47. Link to Digging for the Truth - Carbon Dating – Atlantis ≈ 5:15 (CLICK ON PICTURE)

48. Big Question
How can we experimentally determine the order of a reaction?

49. Make “3” Graphs
In order to determine order of reactant, A. We must collect data consisting of concentration versus time.
“One” common way to determine concentration vs. time data is through the use of a spectrophotometer and graphical information based on this data.

50. Make “3” Graphs We use the data to make three graphs.
[A] versus t
ln [A] versus t
1 / [A] versus t
By examining these graphs we can determine the order of the reaction with respect to a particular reactant and determine the rate constant.
A graphing calculator can be used to make these graphs.

51. [A] versus t (linear for a zero order reaction)
k must be a positive number.

52. ln [A] versus t (linear for a 1st order reaction)

53. 1 / [A] versus t (linear for a 2nd order reaction)

54. Collision Model Key Idea: Molecules must collide to react.
However, only a small fraction of collisions produces a reaction. Why?

55. Two Factors
Collisions must have enough energy to produce the reaction (must equal or exceed the activation energy).
Orientation of reactants must allow formation of new bonds.

56. 2HI → H2 + 2I

57. Concentration and Collision Theory
Why does an increase in concentration cause an increase in reaction rate?

58. Concentration and Collision Theory
Why does an increase in concentration cause an increase in reaction rate?
Increasing the concentration increases the number of collisions and therefore there are more collisions leading to product.

59. Temperature and Collision Theory
Why does a temperature increase cause the reaction rate to increase?

60. Temperature and Collision Theory
Why does a temperature increase cause the reaction rate to increase?
At higher temperatures there are more collisions and a greater percentage of the collisions have the energy necessary to create a successful collision.

61. Activation Energy
The activation energy is the minimum amount of energy necessary for a reaction to occur.

62. Temperature and Activation Energy (Ea)

63. Activation Energy
The activation energy can also be thought of as the energy necessary to form an activated complex during a collision between reactants.

64. Transition State Theory
The activated complex is a hypothetical species lying between reactants and products at a point on the reaction profile called the transition state.

65. The activated complex is a transition state between reactants and products where old bonds have begun to break and new bonds have started to form. It cannot be isolated.

66. For two reactions at the same temperature, the reaction with the higher activation energy has the lower rate constant (k) and the slower rate.

67. 2O3  3O2
A chemical equation like the one above does not tell us how reactants become products - it is simply a summary of the overall reaction.

68. The reaction: 2O3  3O2 O3  O2 + O O3 + O  2O2
Is proposed to occur through the two step process given below:
O3  O2 + O
O3 + O  2O2
This two step process is an example of a reaction mechanism

69. Reaction Mechanisms
A reaction mechanism is a step-by-step description of a chemical reaction.
Each step is called an elementary reaction.

70. Often Used Terms
Intermediate: formed in one step and used up in a subsequent step and so is never seen as a product.
Molecularity: the number of species that must collide to produce the reaction indicated by that step.
Elementary Step: A step within a reaction mechanism whose rate law can be written from its molecularity.

71. Reaction Mechanisms
Elementary Steps
Molecularity: the number of molecules present in an elementary step.
Unimolecular: one molecule in the elementary step.
Bimolecular: two molecules in the elementary step.
Termolecular: three molecules in the elementary step.

72. Reaction Mechanisms Elementary Steps
It is not common to see termolecular processes (statistically improbable).
Unimolecular reactions occur because collisions with other molecules provide the activation energy for the molecule to react.
Bimolecular reactions involve the collision of two particles with sufficient energy and proper orientation.
Termolecular reactions involve the simultaneous collision of three particles with sufficient energy and proper orientation.

73. Reaction Mechanisms Rate Laws for Elementary Steps
The rate law of an elementary step is determined by its molecularity:
Unimolecular processes are first order,
Bimolecular processes are second order, and
Termolecular processes are third order.

74. Reaction Mechanisms
Rate Laws for Elementary Steps

75. The Rate Determining Step

76. Rate-Determining Step
In a reaction mechanism, the rate determining step is the slowest step. It therefore determines the rate of reaction.

77. Reaction Mechanisms Reaction mechanisms must be consistent with:
Stoichiometry for the overall reaction.
The experimentally determined rate law.

78. NO2(g) + CO(g)  NO(g) + CO2(g)
Reaction mechanism must be consistent with the stoichiometry of the overall reaction.
Is the mechanism below consistent with the overall reaction above?
NO2(g) + NO2(g)  NO3(g) + NO(g)
NO3(g) + CO(g)  NO2(g) + CO2(g)

79. Determining the stoichiometry of a reaction mechanism.
Page 439

80. Reaction Mechanisms
The reaction mechanism must also support the rate law.

81. Reaction Mechanisms Rate Laws for Multistep Mechanisms
with an initial fast step.
Consider the reaction:
2NO(g) + Br2(g)  2NOBr(g)

82. Reaction Mechanisms Mechanisms with an Initial Fast Step
2NO(g) + Br2(g)  2NOBr(g)
The experimentally determined rate law is
Rate = k[NO]2[Br2]
Consider the following mechanism

83. The rate law is (based on Step 2):
Rate = k2[NOBr2][NO]
The rate law should not depend on the concentration of an intermediate (intermediates are usually unstable).
NOBr2 is an unstable intermediate, so we express the concentration of NOBr2 in terms of NO and Br2 Since there is an equilibrium in step 1 we have

84. By definition of equilibrium:
Therefore, the overall rate law becomes
Note the final rate law is consistent with the experimentally observed rate law.

85. Student Example: Determine the rate law for the reaction and the balanced equation given the mechanism below:
2NO ↔ N2O2 fast N2O2 + O2 → 2NO2 slow

86. ↔ ↔ Page 439 Assume the rate law is: Rate = k[H2O2][H3O+][I-]
Which step would be the rate – determining step?
Page 439 ↔ ↔

87. This diagram shows a two-step mechanism for a reaction with the first step being rate determining. 

88. What is the mechanism for the reaction?
Overall Reaction

89. Mechanism for Previous Reaction
NO + H2 → NOH2 slow
NO + NOH2 → N2O + H2O fast

90. Catalysts
A catalyst is a substance that increases the rate of a chemical reaction by reducing the activation energy, but which is left unchanged by the reaction.
Catalysis is the process of using a catalyst to speed up a reaction.

92. What is the overall reaction?
O3  O2 + O O3 + O  2O2

93. What is the overall reaction?

94. Identify the intermediates.

95. Identify the intermediates.
NO is a catalyst
A homogeneous catalyst is of the same phase as the reacting substances. It lowers the activation energy by forming intermediates which allow the reaction to proceed by a different pathway.

96. Heterogeneous Catalysts
A heterogeneous catalyst is of a different phase than the reacting substances.
It provides a surface on which the transition state is stabilized thus lowering the activation energy and increasing the reaction rate.
Heterogeneous catalysts are often referred to as surface catalysts.

98. Catalytic Converter: A Heterogeneous Catalyst
In a catalytic converter, the catalyst (in the form of platinum and palladium) is coated onto a ceramic honeycomb that are housed in a muffler-like package attached to the exhaust pipe. The catalyst helps to convert harmful exhaust gases into safer ones.

99. Enzyme Catalysis
Some enzymes accelerate reactions by binding to the reactants in a way that lowers the activation energy.
Other enzymes react with a reactant species to form a new intermediate.
Enzyme catalysis essentially occurs when substances catalyze reactions within a living organism.

100. Substrate Interactions with the Active Sites in Enzyme Catalysis

101. Reaction Rate Lab

102. Reaction Rate Lab – Part A
Use different containers for Reaction Mixtures I and II.
Don’t forget the starch.

104. Clock Reaction Colorless to Blue ≈ 20s

105. Iodine clock reaction ≈ 30s

106. Reaction Rate Lab – Part A
In part A you will perform five different trials with various concentrations.
One side of the table will do Rxn Mix I and the other side will do Rxn Mix II.

107. Reaction Rate Lab – Part B
In part B you will perform trial 1 using a catalyst.