1. Speed or rate of a chemical reaction
Chemical Kinetics
Kinetics
Speed or rate of a chemical reaction

2. Factors that influence reaction rate

3. Measuring Rate

4. Change of Rate with Time C4H9Cl + H2O → C4H9OH + HCl
What happens to the rate of this reaction with respect to time?

5. Change of Rate with Time C4H9Cl + H2O → C4H9OH + HCl
What happens to the rate of this reaction with respect to time?

6. Change of Rate with Time C4H9Cl + H2O → C4H9OH + HCl

7. Rates and Stoichiometry
A → B
If the rate of disappearance of A is 5M/s what is the rate of appearance of B?

8. Rates and Stoichiometry
A → 2B
If the rate of disappearance of A is 5M/s what is the rate of appearance of B?

9. In General aA + bB → cC + dD a t bt ct dt
Rate = [A] = -1[B] = 1[C] = 1[D]
a t bt ct dt

10. Rates and Stoichiometry
A → 2C
If the rate of disappearance of A is 5M/s what is the rate of appearance of C?

11. Concentration and Rate
Increase reactant concentration, increase rate.
But by how much???
Rate Law

12. A → B [A] (M) Rate (M/sec) .1 2x10-3 .2 4x10-3 .4 8x10-3
Determine (a) the rate law, (b) the rate constant, (c) the rate when [A] = .15 M

13. A → B [A] (M) Rate (M/sec) .1 1x10-3 .2 4x10-3 .4 16x10-3
Determine (a) the rate law, (b) the rate constant, (c) the rate when [A] = .3 M

14. A + B → C [A] (M) [B] (M) Rate (M/sec) .1 .1 4x10-3 .1 .2 4x10-3
Determine (a) the rate law, (b) the rate constant, (c) the rate when [A] = .05 M and [B] = .1 M

15. Concentration of Reactants and Time
Rate law expressions allow us to calculate the rate of a reaction.
These rate law expressions can then be converted to equations, using calculus, that can tell us what the concentration of a reactant is at a specific time

16. First Order Reaction
A→B
rate = -[A]
 t

17. First Order Reaction
A→B
rate = -[A] rate = k[A]
 t

18. First Order Reaction
A→B
rate = -[A] rate = k[A]
 t
-[A] = k[A]

19. This is called the differential rate law
First Order Reaction
A→B
rate = -[A] rate = k[A]
 t
-[A] = k[A]
This is called the differential rate law

20. First Order Reaction
ln [A]t - ln [A]o = -kt

21. First Order Reaction
ln [A]t - ln [A]o = -kt
ln [A]t = -kt
[A]o

22. First Order Reaction ln [A]t = -kt ln [A]t - ln [A]o = -kt [A]o
ln [A]t = -kt + ln[A]o

23. ln [A]t = -kt Integrated Rate Laws
First Order Reaction
ln [A]t - ln [A]o = -kt
ln [A]t = -kt Integrated Rate Laws
[A]o
ln [A]t = -kt + ln[A]o

24. The decomposition of dinitrogen pentoxide is a first order reaction with a rate constant of 5.1 x s-1. (a) If the initial concentration was 0.25M, what is the concentration after 3.2 minutes? (b) How long will it take for the concentration to decrease to 0.15M?
2N2O5(g) → 4NO2(g) + O2(g)

25. First Order Reaction
ln [A]t = -kt + ln[A]o

26. First Order Reaction
ln [A]t = -k t + ln[A]o
↓ ↓ ↓ ↓
y = m x + b

27. Second Order Reaction
A →B
rate = -[A]
 t

28. Second Order Reaction
A →B
rate = -[A] rate = k[A]2
 t

29. Second Order Reaction
A →B
rate = -[A] rate = k[A]2
 t
-[A] = k[A]2

30. Second Order Reaction A →B rate = -[A] rate = k[A]2  t
-[A] = k[A]2
This is called the differential rate law

31. Second Order Reaction
Integrated Rate Law

32. Half-Life (t1/2 )
Time required for the concentration of the reactant to reach ½ of its initial value.

33. Half-Life (t1/2 )
Time required for the concentration of the reactant to reach ½ of its initial value.
[A]t ½ = ½[A]o

34. Half-Life (t1/2 )
Time required for the concentration of the reactant to reach ½ of its initial value.
[A]t ½ = ½[A]o
First order reaction : t1/2 = /k

35. Half-Life (t1/2 )
Time required for the concentration of the reactant to reach ½ of its initial value.
[A]t ½ = ½[A]o
First order reaction : t1/2 = /k
Second order reaction : t1/2 = 1/k[A]o

36. Note: nuclear decay is a ? order reaction
Half-Life (t1/2 )
Time required for the concentration of the reactant to reach ½ of its initial value.
[A]t ½ = ½[A]o
First order reaction : t1/2 = /k
Second order reaction : t1/2 = 1/k[A]o
Memorize!
Note: nuclear decay is a ? order reaction

37. The recombination of iodine atoms to form molecular iodine in the gas phase follows second order kinetics and has a rate constant of 7.0 x 109 M-1. (a) If the initial concentration of I was 0.086M, calculate the concentration after 2.0 minutes. (b) Calculate the half-life of the reaction if the initial concentration of I is 0.60M.
I + I → I2

38. In order for a reaction to occur, particles must collide
Collision Theory
In order for a reaction to occur, particles must collide

39. The particles must collide at the proper angle
Two conditions
1. Orientation
The particles must collide at the proper angle

40. Two conditions
2. Energy

41. The particles must collide with enough energy to stick
Two conditions
2. Energy
The particles must collide with enough energy to stick

42. Cons
Ran

43. Arrhenius Equation
Shows the quantitative relationship between rate, activation energy, and temperature

44. Calculations Involving Ea
k = Ae –Ea/RT

45. Calculations Involving Ea
k = Ae –Ea/RT
ln k = ln Ae–Ea/RT

46. Calculations Involving Ea
k = Ae –Ea/RT
ln k = ln Ae–Ea/RT
ln k = ln A – Ea/RT

47. Calculations Involving Ea
k = Ae –Ea/RT
ln k = ln Ae–Ea/RT
ln k = ln A – Ea/RT
ln k = (-Ea/R) (1/T) + ln A

48. Calculations Involving Ea
k = Ae –Ea/RT
ln k = ln Ae–Ea/RT
ln k = ln A – Ea/RT
ln k = (-Ea/R) (1/T) + ln A
y = m x b

49. ln k = (-Ea/R) (1/T) + ln A
y = m x b

50. Calculations Involving Ea

51. Reaction Mechanisms
Most chemical reactions occur in a series of steps, not a single one as shown by a balanced chemical equation

52. Reaction Mechanisms
Most chemical reactions occur in a series of steps, not a single one as shown by a balanced chemical equation
The sequence of steps a reaction goes through is called the reaction mechanism

53. Reaction Mechanisms Elementary Reactions
Reactions that do occur in a single step

54. Reaction Mechanisms Elementary Reactions
Reactions that do occur in a single step
NO + O3 → NO2 + O2

55. Reaction Mechanisms Elementary Reactions
Reactions that do occur in a single step
NO + O3 → NO2 + O2

56. Reaction Mechanisms Multistep Reactions
NO2 + CO → NO + CO2

57. Reaction Mechanisms Multistep Reactions
NO2 + CO → NO + CO2
Occurs in 2 steps

58. Reaction Mechanisms Multistep Reactions
NO2 + CO → NO + CO2
Occurs in 2 steps
NO2 + NO2 → NO3 + NO

59. Reaction Mechanisms Multistep Reactions
NO2 + CO → NO + CO2
Occurs in 2 steps
NO2 + NO2 → NO3 + NO
NO3 + CO → NO2 + CO2

60. Rate Laws (elementary reactions)
For an elementary reaction and only an elementary reaction we can use the coefficients from the balanced equation as the reaction orders.

61. Table 14.03 Figure 14-T03 Title: Table 14.3 Caption:
Elementary Reactions and Their Rate Laws
Notes:
Keywords:

62. Rate Laws (multistep reactions)
The slowest reaction in a multistep reaction will determine the rate of the reaction

63. Rate Laws (multistep reactions)
Reactions where the first step is the slowest

64. Rate Laws (multistep reactions)
Reactions where the first step is the slowest
NO2 + CO → NO + CO2
NO2 + NO2 → NO3 + NO (slow)
NO3 + CO → NO2 + CO2 (fast)

65. Rate Laws (multistep reactions)
Reactions where the second step is the slowest

66. Rate Laws (multistep reactions)
Reactions where the second step is the slowest
2NO + Br2 → 2NOBr
NO + Br2 ⇌ NOBr2 (fast)
NOBr2 + NO → 2NOBr (slow)

67. Catalyst Catalyst increase rate without being used up.
They are very specific
They work by lowering activation energy

68. Catalyst work by lowering activation energy

69. Catalyst
Homogeneous
Same phase as reacting molecules

70. Catalyst Heterogeneous Different phase as reacting molecules
Work by adsorption

71. Catalyst
Enzymes

72. Catalyst
Enzymes
Large molecular mass proteins

73. Lock and Key Model