1. Chemical Kinetics
Chapter 12

2. Objectives Define key terms and concepts.
Predict how temperature, catalysts, concentration, and surface area affect reaction rates.
Calculate the rate of a reaction.
Distinguish between a zero, first, and second order reactions.
Identify the rate determining step of a reaction.

3. Rates of Reaction Chemical Kinetics Reaction Rate
Area of chemistry that studies the rate of reactions.
Reaction Rate
The speed a chemical reaction occurs.
Rate of a reaction may change over time.
May also have a back-reaction occurring later in the reaction that wasn’t occurring at the start.

4. Rates of Reaction Rate Constant (k) Rate Law
A constant the relates the reaction rate and the concentration of the reactants.
Holds true for the entire reaction.
Rate Law
Expression relating the rate of a reaction to the rate constant and the concentrations of the reactants.

5. Pb(NO3)2(aq) + 2KI(aq)  2KNO3(aq) + PbI2(s)
Rate of Reaction
Lead Iodide
Pb(NO3)2(aq) + 2KI(aq)  2KNO3(aq) + PbI2(s)
Iodine Clock
Reaction of hydrogen peroxide and sulfuric acid with potassium iodide, sodium thiosulfate, and starch.

6. Affects on Rates of Reaction
Collision Theory of Chemical Kinetics
The rate of reaction is directly proportionate to the number of collisions per second.
Temperature
Collision Orientation
Surface Area

7. Zn(s) + HCl(aq)  ZnCl2(aq) + H2(g)
Rates of Reaction
Concentration
Increasing/decreasing the concentration of either of the reactants will also increase the reaction rate proportionately.
If the reaction order is 0 for a reactant, the rate is independent of the concentration of that reactant.
In the lab, zinc granules react fairly slowly with dilute hydrochloric acid, but much faster if the acid is concentrated.
Zn(s) + HCl(aq)  ZnCl2(aq) + H2(g)

8. Affects on Rates of Reaction
Activation Energy (Ea)
The minimum amount of energy required to initiate a chemical reaction.
If the products of a reaction are more stable than the reactants, the reaction is exothermic.
If the products of a reaction are less stable than the reactants, the reaction is endothermic.

9. Affects on Rates of Reaction

10. Rates of Reaction
The Reaction Rate is expressed in terms of the decrease in the concentration of reactants as compared to the increase in the concentration of products over time.
For the general reaction
aA  bB
Where [A] is the change in concentration of the reactants
[B] is the change in concentration of the products
t is time in seconds 1 a
∆[A] ∆t 1 b
∆[B] ∆t
Rate = -
Rate =

11. Rates of Reaction For the general reaction aA + bB  cC + dD 1 a ∆[A]

12. If aluminum is reacting at the rate of 0
If aluminum is reacting at the rate of 0.015M/second with chlorine gas as illustrated in the following reaction, Al + Cl2  AlCl3 at what rate is aluminum chloride being formed?

13. If aluminum is reacting at the rate of 0
If aluminum is reacting at the rate of 0.015M/second with chlorine gas as illustrated in the following reaction, Al + Cl2  AlCl3 at what rate is chlorine gas being consumed?

14. If silver nitrate reacts at a rate of 0
If silver nitrate reacts at a rate of 0.094M/sec in the following reaction AgNO3 + CaCl2  Ca(NO3)2 + AgCl at what rate is calcium chloride being consumed? At what rate are calcium nitrate and silver chloride being produced?

15. What Are Your Questions?

16. Rates of Reaction For the general reaction aA + bB  cC + dD
Reaction Order
The sum of the powers to which all reactant concentrations appear in the rate law are raised.
Given by adding together x and y in the above reaction.
If for a reaction, x = 1 and y = 2, then overall the reaction order is 3.
Rate = k[A]x[B]y

17. Rates of Reactions

18. Rates of Reaction – First Order
The reaction rate and concentration of reactants can be calculated for a reaction at any time during that reaction.
First Order Reactions
Reaction whose rate depends on the reactant concentration raised to the first power.
Rate=k[A]
[A]t
[A]0
In = -kt
In[A]t = -kt + ln[A]0

19. Rates of Reaction – First Order

20. The following reaction is first order with a rate constant of 4
The following reaction is first order with a rate constant of 4.80x10-4/sec at 45°C. N2O5  NO2 + O2 If the initial concentration is 1.65x10-2M, what is the concentration after 825sec? How long would it take for the concentration of N2O5 to decrease to 1.00x10-2M from the initial value provided?

21. The following reaction is first order with a rate constant of 4
The following reaction is first order with a rate constant of 4.80x10-4/sec at 45°C. N2O5  NO2 + O2 If the initial concentration is 1.65x10-2M, what is the concentration after 600 seconds? How long would it take for the concentration of dinitrogen pentoxide to decrease to 10% of its initial value?

22. Rates of Reaction – Half Life
The amount of time it takes for the concentration of a reactant to decrease by half.
Radioactivity
t1/2 = 0.693 k

23. Sulfuryl chloride, SO2Cl2, is a colorless, corrosive liquid whose vapor decomposes in a first- order reaction to sulfur dioxide and chlorine. SO2Cl2(g)  SO2(g) + Cl2(g) At 320°C, the rate constant is 2.20x10-5/sec. What is the half-life of SO2Cl2 at this temperature? How long would it take for 50% of the SO2Cl2 to decompose? How long for 75%?

24. Rates of Reaction – Second Order
Second Order Reaction
A reaction in which the rate depends on the concentration of 1 reactant raised to the 2nd power or 2 different reactants, raised to the 1st power.
Rate=k[A]2 1 1
t1/2 = 1 =
+ kt
[A]t
[A]0
k[A]0

25. Rates of Reaction – Second Order

26. Rates of Reaction – Zero Order
Zero Order Reaction
The concentration of the reactants does not affect the rate of the reaction.
Rate = k
[A]t = - kt + [A]0
t1/2 = [A]0 2k

27. Rates of Reaction – Zero Order

28. Summary of the Rate Laws

29. Hydrogen sulfide is oxidized by chlorine in aqueous solution according to the following reaction H2S(aq) + Cl2(aq)  S(s) + 2HCl(aq) If the rate law is experimentally determined as Rate=k[H2S][Cl2] What is the reaction order with respect to each of the reactants? What is the overall reaction order?

30. Oxalic acid, H2C2O4, is oxidized by permanganate in aqueous solution according to the following reaction 2MnO4-(aq) + 5H2C2O4(aq) + 6H+(aq) 2Mn2+(aq) + 10CO2(aq) + 8H2O(l) If the rate law is experimentally determined as Rate=k[MnO4-][H2C2O4] What is the reaction order with respect to each of the reactants? What is the overall reaction order?

31. A reaction is zero order with a rate constant of 5. 4x108/M•sec
A reaction is zero order with a rate constant of 5.4x108/M•sec. If the amount of the reactant is increased, what is the rate of reaction? What is the half-life of this reaction if the concentration of the products is 0.25M?

32. Using the following reaction and initial rates of reaction provided, determine the reaction order for the reactants and the rate constant for the reaction. 2NO2(g)  2NO(g) + O2(g)
Experimental Trial
Initial NO2 Concentration (M)
Initial Rate of Formation of O2 (mole/L*sec)
Exp 1
0.010
7.1x10-5
Exp 2
0.020
28x10-5

33. Initial Concentration (M) Initial Rate (mole/L*sec)
Using the following reaction and initial rates of reaction provided, determine the reaction order for the reactants and the rate constant for the reaction. H2O2(aq) + 3I-(aq) + 2H+(aq)  I3-(aq) + 2H2O(l)
Experimental Trial
Initial Concentration (M)
Initial Rate (mole/L*sec)
H2O2 I- H+
Exp 1
0.010
1.15x10-6
Exp 2
0.020
2.30x10-6
Exp 3
Exp 4

34. HO2 is a highly reactive chemical that can form in the atmosphere
HO2 is a highly reactive chemical that can form in the atmosphere. It can decompose into hydrogen peroxide and oxygen according to the following reaction: HO2(g)  H2O2(g) + O2(g) Determine the reaction order and rate constant at 25C using the following data:
Trial
[HO2]0 (M)
Initial Rate (M/s) 1
1.1x10-8
1.7x10-7 2
2.5x10-8
8.8x10-7 3
3.4x10-8
1.6x10-6 4
5.0x10-8
3.5x10-6

35. What Are Your Questions?

36. The Arrhenius Equation
Expresses how temperature effects the rate constant of a reaction
Where Ea is the activation energy in kJ/moles
R is J/K•mole
T is the temperature in kelvins
k = Ae-Ea/RT
lnk = lnA - Ea/RT

37. The rate constant for the formation of hydrogen iodide is 2
The rate constant for the formation of hydrogen iodide is 2.7x10-4L/mole•sec at 600K and 3.5x10-3L/mole•sec at 650K. What is the activation energy for this reaction? What is the rate constant at 700K? H2(g) + I2(g)  2HI(g)

38. Acetaldehyde decomposes when heated as illustrated in the following reaction: CH3CHO(g)  CH4(g) + CO(g) If the rate constant for this decomposition reaction is 1.05x10-3L/mole•sec at 760K and 2.14x10-2L/mole•sec at 830K. What is the activation energy for this reaction? What is the rate constant at 880K?

39. What Are Your Questions?

40. Reaction Mechanisms
A balanced chemical reaction only explains the overall reaction process, but not all the individual steps it required to make the reaction happen.
Elementary Steps
Series of simple reactions that represent the progress of an overall reaction at the molecular level.
Reaction Mechanism
The sequence of elementary steps that lead to the products of a reaction.

41. Reaction Mechanisms CH2=CH2 + HCl  CH3CH3Cl Intermediates
A compound that appears in the elementary steps of the reaction, but not the overall reaction.

42. Reaction Mechanisms
Need to be aware of the number of molecules reacting in an elementary step to determine the rate order.
Bimolecular Reaction
Involves 2 molecules
Unimolecular Reaction
Involves 1 molecule
Decomposition or dehydration reactions
Termolecular Reactions
Involves 3 molecules

43. Reaction Mechanisms Rate-Determining Step
The slowest in the sequence of steps in a chemical reaction that leads to the formation of the products.
Back to the Iodine Clock
In the first, slow reaction, the triiodide ion is produced
H2O2(aq) + 3 I−(aq) + 2 H+(aq) → I3−(aq) + 2 H2O (l)
In the second, fast reaction, triiodide is reconverted to iodide by the thiosulfate.
I3−(aq) + 2 S2O32−(aq) → 3 I−(aq) + S4O62-(aq)

44. Reaction Mechanisms Overall Balanced Reaction
2HC2H3O2(aq) + Na2CO3(aq)  2C2H3O2Na(aq) + H2O(l) + CO2(g)
Step 1
2HC2H3O2(aq) + Na2CO3(aq)  2C2H3O2Na(aq) + H2CO3(aq)
Step 2
H2CO3(aq)  H2O(l) + CO2(g)

45. Reaction Mechanisms When determining the Rate Law
The sum of the elementary steps must give the overall balanced equation for the reaction.
The rate-determining step should show the same rate law as predicted experimentally.

46. Carbon tetrachloride is produced by chlorinating methane or a chlorinated methane molecule. The mechanism for the overall reaction is provided below: Cl2 2Cl Cl + CHCl3 HCl + CCl3 Cl + CCl3 CCl4 What is the overall, balanced net reaction? Identify the intermediates.

47. O3(g) + 2NO2(g)  O2(g) + N2O5(g)
Ozone reacts with nitrogen dioxide according to the following equation:
O3(g) + 2NO2(g)  O2(g) + N2O5(g)
The proposed steps for this reaction are
O3(g) + NO2(g)  NO3(g) + O2(g) (slow step)
NO3(g) + NO2(g)  N2O5(g) (fast step)
What is the rate law predicted by this mechanism?

48. H2O2(aq) + I-(aq)  H2O(l) + IO-(aq) (slow step)
The iodide ion-catalyzed decomposition of hydrogen peroxide occurs by the following mechanism:
H2O2(aq) + I-(aq)  H2O(l) + IO-(aq) (slow step)
H2O2(aq) + IO-(aq)  H2O(l) + O2(g) + I-(aq) (fast step)
What is the rate law predicted by this mechanism?

49. Catalysis
Catalyst
Something that speeds up the rate of a reaction by lower the activation energy required for the reaction.
Is not used up in the reaction.

50. Catalysis Heterogeneous Catalysis Homogeneous Catalysis
The reactants and the catalyst are in different phases.
Synthesis of Ammonia
Fritz Haber
Catalytic Converters
Homogeneous Catalysis
When the reactants and catalyst are in the same phase, typically liquid.

51. Catalysis

52. What Are Your Questions?