1. Chemical Kinetics
SCH4U: Grade 12 Chemistry

2. Unit Mind Map

3. Chemical Kinetics Chemical kinetics? Why do we study this?
Who makes use of chemical kinetics?

4. Chemical Kinetics
Chemical kinetics  studying how we can make reactions go faster or slower
Why do we study this?
Who makes use of chemical kinetics?

5. Chemical Kinetics
Chemical kinetics  studying how we can make reactions go faster or slower
Why should we study this?
Economics effects Materials are expensive
Developing pharmaceutical drugs
Who makes use of chemical kinetics?

6. Chemical Kinetics
Chemical kinetics  studying how we can make reactions go faster or slower
Why should we study this?
Economics effects Materials are expensive
Developing pharmaceutical drugs
Who makes use of chemical kinetics?
Biologists  metabolic rxns, food digestion, bone regeneration
Automobile engineers rate of rusting, decrease pollutants
Agriculture  slow down food ripening

7. Reaction Rates What is a chemical reaction rate?
Possible formula for measuring rate?

8. Reaction Rates What is a chemical reaction rate?
measure of how fast reactants are used up or how fast products are produced
Possible formula for measuring rate?

9. Reaction Rates What is the reaction rate?
measure of how fast reactants are used up or how fast products are produced
Possible formula for measuring rate?
Rate = Δ Concentration/Δ Time

10. Sample problem N2 (g) + 3H2 (g)  2NH3(g)
What is the average rate of production of ammonia for the system if the concentration is 3.5mol/L after 1.0min and 6.2mol/L after 4.0 min?

11. Sample problem N2 (g) + 3H2 (g)  2NH3(g)
What is the average rate of production of ammonia for the system if the concentration is 3.5mol/L after 1.0min and 6.2mol/L after 4.0 min?
ANSWER: 0.9mol/(L*min)

12. Measuring reaction rates Graphically
Average rate of reaction  take the slope of the secant of the line
Instantaneous rate of reaction  take the slope of the tangent of the line

13. Measuring reaction rates Graphically
Average rate of reaction  take the slope of the secant of the line
Draw a secant line between two points
Calculate the slope
Slope = Rise/Run
= Δconcentration/Δ time

14. Measuring reaction rates Graphically
Instantaneous rate of reaction  take the slope of the tangent of the line
Draw a tangent line to the graph
Calculate the slope of the tangent line
Slope = Rise/Run
= Δconcentration/Δ time

15. Measuring reaction rates
What are some factors that we can measure experimentally to determine the rate of reaction ?

16. Measuring reaction rates
Production of a gas
Production of ions
Changes in colour

17. Reaction rates and stoichiometry
CO(g) + NO2(g)  CO2(g) + NO(g) In this reaction, the ratio of CO to NO is 1:1 Therefore, the disappearance of CO is the same as the production of NO Rate = - Δ [CO]/Δt = + Δ[NO]/Δt

18. Reaction rates and stoichiometry
Suppose the ratio is NOT 1:1? Example, H2 (g) + I2 (g)  2HI (g) 2 mols of HI are produced for every 1 mol H2 used Rate = The rate at which H2 is used up is only half of which HI is produced

19. Sample Problem IO3- (aq) + 5I- (aq) + 6H+ (aq)  3I2 (aq) + 3H2O (l)
The rate of consumption of iodate ions (IO3-) is determined experimentally to be 3.0 x 10-5 mol/(L*s). What are the rates of reaction or the other reactants and products in this reaction?
Complete p.364 #3, 4
Homework: practice problem package

20. Collision Theory
Why do reactions occur the way that they do? VS

21. Collision Theory Collision theory  reactions can only occur if:
There is a collision between molecules

22. Collision Theory Collision theory  reactions can only occur if:
1. There is a collision between molecules
2. the molecules are oriented in the correct way

23. Collision Theory Collision theory  reactions can only occur if:
1. There is a collision between molecules
2. the molecules are oriented in the correct way
3. enough energy is provided to break the chemical bonds that hold molecules together (Activation energy)

24. Activation Energy
Activation Energy  minimum potential energy the system needs to overcome for the molecules to react

25. Factors That Affect Rate of Reaction
What are some of the factors that may speed up or slow down a chemical reaction? P.392

26. Factors That Affect Rate of Reaction
What are some of the factors that may speed up or slow down a chemical reaction?
Chemical nature of the reactant
Concentration of reactants
Surface area of reactants
Temperature
Catalysts
Get into groups, brainstorm and then explain these factors to the rest of class in a creative way, making use of collision theory

27. Lesson 2 Review collision theory Rate Laws Group Practice problems
Individual practice problems
homework

28. Unit Mind Map

29. Collision theory

30. Collision Theory Collision theory  reactions can only occur if:
1. There is a collision between molecules
2. the molecules are oriented in the correct way
3. enough energy is provided to break the chemical bonds that hold molecules together (Activation energy)

31. Factors That Affect Rate of Reaction

32. Factors That Affect Rate of Reaction
Chemical nature of the reactant
Concentration of reactants
Surface area of reactants
Temperature
Catalysts

33. The Rate Law
Mathematical relationship between reaction rate and factors that affect it
Determined empirically (experimentally)
Rate = k[X]m[Y]n
e.g. 2NO2 + F2  2NO2F
The above reaction is 1st order with respect to NO2 and 2nd order with respect to F2. What is the rate law equation?

34. The Rate Law
Answer:
e.g. 2NO2 + F2  2NO2F
rate = k[NO2]1[F2]2

35. Steps to solve rate law problems
eg. 2BrO3- (aq) + 5HSO3- (aq)  Br2 (g) + 5SO42- (aq) + H2O (l) + 3H+ (aq)
Rate = k [BrO3-]m [HSO3-]n

36. Steps to solve rate law problems
eg. 2BrO3- (aq) + 5HSO3- (aq)  Br2 (g) + 5SO42- (aq) + H2O (l) + 3H+ (aq)
Write out rate equation: rate = k [BrO3-]m [HSO3-]n
pick a trial where [BrO3-] changes but [HSO3-] stays constant= trial 1 and trial 2
Using a ratio, determine the relationship between change in concentration and change in rate = (trial 1/trial2) = (4.0/2.0) = (1.6/0.8)
= 2 = 2
4. Using the following chart, determine the rate order of [BrO3-]

37. Steps to solve rate law problems
5. If the concentration is doubled (2.0 to 4.0), the rate also doubles (0.80 to 1.60), therefore the rate order is 1 m=1

38. Steps to solve rate law problems
6. To find the rate order of the other reactant, follow the same steps but pick a trial where [HSO3-] changes but [BrO3-] stays constant= trial 2 and trial 3 = (6.0/3.0) = (0.8/0.2) = 2 = 4 7. Use the following table to determine the rate order

39. Steps to solve rate law problems
8. As the concentration doubles the rate was multiplied by 4, therefore rate order of HSO3- = 2 (n=2) 9. Plug values into rate equation rate = k [BrO3-]m [HSO3-]n and solve for k.

40. Example 1
From the data collected above, determine the rate law of the following equation:
aA + bB  products
Rate = k[A]m[B]n

41. Example 1 aA + bB  products Rate = k[A]m[B]n m = 0; n = 1; k = 0.6s-1
Rate = 0.6s-1 [B]1

42. Example 2
From the data collected above, determine the rate law of the following equation:
2NO2 + F2  2NO2F
Rate = k[A]m[B]n

43. Example 2 2NO2 + F2  2NO2F Rate = k[A]m[B]n
m = 2, n = 0, k = 4 x 10-3M-1s-1
Rate =4 x 10-3M-1s-1 [NO2]2

44. Team problem Get into groups of 3 and obtain problem clues
a) Determine the order with respect to each reactant
b) Determine the overall order of reaction c) Write the rate expression for the reaction. d) Find the value of the rate constant, k.

45. Practice Problem
From the data collected above, determine the rate law for the following equation:
IO3- (aq) + 5I- (aq) + 6H+ (aq)  3I2 (aq) + 3H2O (l)

46. Practice Problem
From the data collected above, determine the rate law for the following equation:
IO3- (aq) + 5I- (aq) + 6H+ (aq)  3I2 (aq) + 3H2O (l)
m = 1; n = 1; o = 2; k = 5M-3s-1
Rate = 5M-3s-1 [IO3-]1 [I-]1 [H+]2

47. Homework
P. 377 practice # 1, 2, 3, 6,
P. 415 # 15 a-e

48. Lesson 3 Reaction mechanisms: Elephant toothpaste demonstration
Ping pong activity
Assembly line
Elephant toothpaste demonstration
Practice problems

49. Unit Mind Map

50. Reaction mechanisms Rate Laws  determined experimentally
Equation: Reactants products
This actually occurs in a series of steps called elementary steps
Analogy: Cooking takes place in several steps

51. Reaction mechanisms
What are the chances of the following reaction occurring in one step?
Hint: collision theory
4HBr (g) + O2 (g)  2H2O (g) + 2Br2 (g) actually occurs in 3 separate steps

52. Reaction mechanisms
4HBr (g) + O2 (g)  2H2O (g) + 2Br2 (g) actually occurs in 3 separate steps:
HBr (g) + O2 (g)  HOOBr (g) (slow)
HOOBr (g) + HBr (g)  2HOBr (g) (fast)
2HOBr (g) + 2HBr (g)  2H2O (g) + 2Br2 (g) (fast)
4HBr (g) + O2 (g)  2H2O (g) + 2Br2

53. Reaction mechanisms
4HBr (g) + O2 (g)  2H2O (g) + 2Br2 (g) actually occurs in 3 separate steps:
HBr (g) + O2 (g)  HOOBr (g) (slow)
HOOBr (g) + HBr (g)  2HOBr (g) (fast)
2HOBr (g) + 2HBr (g)  2H2O (g) + 2Br2 (g) (fast)
4HBr (g) + O2 (g)  2H2O (g) + 2Br2
Rate law: rate = k[HBr2][O2]

54. Reaction mechanisms: working backwards
Overall equation: 4HBr (g) + O2 (g)  2H2O (g) + 2Br2 (g) Rate law: rate = k[HBr]1[O2]1 Devise a proposed mechanism for this reaction

55. Reaction mechanisms: working backwards
Overall equation: 4HBr (g) + O2 (g)  2H2O (g) + 2Br2 (g) Rate law: rate = k[HBr][O2] Devise a proposed mechanisms for this reaction Step 1. using the rate law, write out the rate- determining (slow) step Step 2. use the overall equation to determine what needs to be added to achieve the overall equation Step 3. cross out intermediates and add up what is left to produce the overall equation

56. Reaction mechanisms: working backwards
4HBr (g) + O2 (g)  2H2O (g) + 2Br2 (g) actually occurs in 3 separate steps:
HBr (g) + O2 (g)  HOOBr (g) (slow)
HOOBr (g) + HBr (g)  2HOBr (g) (fast)
2HOBr (g) + 2HBr (g)  2H2O (g) + 2Br2 (g) (fast)
4HBr (g) + O2 (g)  2H2O (g) + 2Br2

57. Potential Energy Diagram

58. Elephant toothpaste demonstration

59. Elephant toothpaste demonstration
Elephant toothpaste demonstration reaction: 2H2O2 O2 + 2H2O Rate = k[H2O2]1[I-]1

60. Elephant toothpaste demonstration
Elephant toothpaste demonstration reaction: 2H2O2 O2 + 2H2O Rate = k[H2O2]1[I-]1 Proposed Reaction Mechanism: H2O2 + I-  IO- + H2O (Slow) H2O2 + IO-  I- + H2O + O2 (Fast) 2H2O2  O2 + 2H2O

61. Elephant toothpaste demonstration
Possible energy potential diagram?

62. Practice Problem
Propose a possible mechanism for the following reaction: 2N2O5 (g)  2N2O4 (g) + O2 (g) r = k[N2O5]1

63. Practice Problem
Propose a possible mechanism for the following reaction: 2N2O5 (g)  2N2O4 (g) + O2 (g) r = k[N2O5]1 Possible mechanism: N2O5  N2O4 + O (Slow) O + N2O5  N2O4 + O2 (Fast) 2N2O5  2N2O4 + O2

64. Practice Problem
Propose a possible mechanism for the following reaction: 2NO2 + F2  2NO2F r = k[NO2]1[F2]1

65. Practice Problem
Propose a possible mechanism for the following reaction: 2NO2 + F2  2NO2F r = k[NO2]1[F2]1 Possible mechanism: NO2 + F2  NO2F + F (Slow) F + NO2  NO2F (Fast)

66. Homework
p. 390 #2 p. 391 # 1, 2, 3 Also, complete worksheet.