1. Kinetics Chemistry—Introduction

2. Kinetics Study of reaction rates Rate
How fast does it happen? What variables influence the rate? What is the path the reaction takes to convert reactants to products?
Rate
Change in an amount over a period of time
Ex. Distance traveled, space travel

4. Reaction Rates
Change in the concentration of a chemical compound in the reaction over a period of time
Focus on one reactant or one product in the reaction
Want to know rate of disappearance for reactant/rate of appearance for a product
Rate = ∆X ∆t
Units = M/time

6. Example 1: A + B → D + E A) Rate of disappearance for A ∆t
RateA = -∆[A] ∆t
B) Rate of appearance for D
RateD = ∆[D]

7. Example 2: 2AB → A2 + B2 Determine the rate of disappearance of AB
Time = 0 seconds
15.0 M 0M
Time= 60 seconds
5.0 M
5.00M
Determine the rate of disappearance of AB
Determine the rate of appearance of A2 and B2

9. Reaction Rate at ANY moment in time
Slope of line tangent to the reaction curve

10. Collision Theory
Reactant molecules MUST collide to produce a chemical reaction
The concentrations of reactants affect the # of collisions among reactants
For reactions occurring in one step—rate of reaction is proportional to product of reactant concentrations
Rate = k[A] [B]
Rate of any reaction step dependent on collision frequency

11. Variables Affecting Reaction Rate
Collision Rates between reactants
% of collisions with reactants arranged in proper orientation to produce reaction.
% of collisions with energy energy (activation energy) to produce reaction.

12. When do collision rates increase?
Increase in concentrations of reactants
Temperature increases
WHY?

13. Most collisions do NOT result in a chemical reaction!
Small percentage of collisions actually convert reactants to products. Why?
Molecular Orientation
Random orientation
Not all collisions have correct orientation
Molecular Energy at Collision
Molecules have different kinetic energies
Collision energy is energy source to get a reaction started

14. Activation Energy (Ea)
The amount of collision energy needed to overcome Ea so the reaction can occur
Amount of energy needed for a chemical reaction to happen, energy needed to convert reactants to products.

15. Activation Energy--Endothermic

16. When will reactions occur?
MUST have a collision
Collision must happen with the correct molecular orientation to generate a reaction
Collision energy ≥ Ea

17. Arrhenius Equation
Rate constant and reaction rate are temperature dependent.
Enables the activation energy for a reaction to be determined based on the relationship between reaction rate and temperature.

18. lnk = -Ea ( 1/T ) + lnA R Arrhenius Equation k = rate constant
Ea = activation energy (J)
R = J/molK
T = Kelvin
Z = proportionality constant, changes based on reaction

19. ln (k1/k2) = Ea (1/T2 – 1/T1) R Arrhenius Equation
Different form of equation can be used to observe how temperature changes affect the rate constant (k)
ln (k1/k2) = Ea (1/T2 – 1/T1) R

20. Example 1
Calculate activation energy (Ea) for HI decomposition with the following data.
Temperature (K)
Rate Constant (M/s)
573
2.91 x 10-6
673
8.38 x 10-4
773
7.65 x 10-2

21. Collision Theory
In order for two particles to react chemically, they must collide. Not only must they collide, but it must be an “effective collision.” That is, they must have the correct amount of energy and collide with the proper orientation in space.
Any factor which increases the likelihood that they will collide will increase the rate of the chemical reaction.
Presented by Mark Langella, PWISTA.com

22. FACTORS WHICH AFFECT THE RATE OF A CHEMICAL REACTION (Bonds Must Break)
Surface Area/ Contact Area (Opportunity for collisions)
Concentration ( Increase Frequency)
Temperature ( Increase Frequency)
Catalyst ( Effective Collisions)
Nature of Reactants ( Effective collisions)
Presented by Mark Langella, PWISTA.com

23. What can influence reaction rates?
Temperature
Concentration
Catalyst
Surface Area
Volume/Pressure
Reactant Properties

24. Plot showing the number of collisions with a particular
energy at T1& T2, where T2 > T1 -- Boltzman Distribution.
Presented by Mark Langella, PWISTA.com

25. Rate Law (cont.) A + B → C + D Rate = k [A]m[B]n
Rate = rate of disappearance of reactants
k = rate constant, specific to reactions and temperature
m = reaction order in terms of A
n = reaction order in terms of B
m + n = overall reaction order

26. Reaction Order
Indicates how concentration changes affect changes in the reaction rate
Orders: 0, 1, 2
Overall order of reaction = Σ individual orders of each reactant
Order of a reaction in terms of a reactant ≠ reactant’s coefficient in chemical equation

27. Reaction Orders (cont.)
Zero-order Reaction
Rate not dependent on reactant’s concentration
Constant reaction rate
First-order Reaction
Concentrate affects reaction rate
Example: Double concentration, double the rate.
Second-order Reaction
Concentration affects reaction rate
Example: double concentration, quadruple the rate

28. Rate Constant (k) Units
Reaction Order
Basic Formula
Units
Rate = k
Ms-1 1
Rate = k [A]
s-1 2
Rate = k [A]2
M-1s-1 3
Rate = k [A]3
M-2s-1

29. Example 3: 2NO(g) + O2(g)  2NO2(g)
Based on the reaction’s rate law of
Rate = k(NO)2 (O2)
Classify this reaction’s order.
Overall reaction order = 3rd order

30. Example 4:
Determine the rate law, reaction order, and rate constant (k) for the following reaction at a specific temperature----
2NO(g) H2(g)  N2(g) H2O(g)
Experiment
[NO]initial
[H2]initial
Rate initial 1
0.20M
0.30M
M/s 2
0.10M
M/s 3
M/s
Rate = k[NO]2[H2], overall reaction order is 3rd order
k = 7.5 M-2s-1

31. Example 5: Determine the rate law for the following reaction----
NH4+(aq) + NO2-(aq)  N2(g) H2O(l)
Experiment
[NH4+]initial
[NO2-]initial
Rate initial 1
5 x 10-2 M
2 x 10-2 M
2.70 x 10-7 M/s 2
4 x 10-2 M
5.40 x 10-7 M/s 3
1 x 10-1 M
Rate law = k(NH4+)1(NO2-)1
Overall reaction order = 2

32. Integrated Rate Law
Enables the determination a reactant’s concentration at any moment in time
Enables the determination of the time it takes to reach a certain reactant concentration
Enables the determination of the rate constant or reaction order

33. 1st Order Integrated Rate Law
Only used with 1st order reactions
Focus on initial concentration and ΔC for one reactant
Initial concentration of reactant known---- can determine reactant concentration at any time
Initial and final reactant concentrations known---can determine rate constant

34. 1st Order Integrated Rate Law
Rate = -Δ[A] = k [A] Δt
-take equation and integrate with calculus to get….
ln[A]t – ln[A]0 = - kt
[A]0 = initial concentration (t = 0)
[A]t = concentration after a period of time

35. Example 1: A  B + 2D
Using the data provided for a 1st order reaction, determine the rate constant and [A] at time = 5.0 x 102s.
Time (s)
[A] (M)
0.020
5.0 x 10
0.017
1.0 x 102
0.014
1.5 x 102
0.012
2.0 x 102
0.010
k = 3.3 x 10-3s-1
[A] = 3.8 x 10-3 M

36. Half-life Radioactive decay is a 1st order process Half-life (t1/2)—
Time it takes for half of a chemical compound to decay or turn into products
Focus on reactant
Constant, not dependent on [ ]
Rate changes with temperature so half-life varies based on temperature

37. Example 2:
Find the half-life for the following reaction with a reate constant (k) of 1.70 x 10-3 s-1
Half life = 408s

38. 2nd Order Integrated Rate Law
Used only for second order reactions
Focus on initial concentration and ΔC for one reactant with reaction 2nd order with respect to it.
Initial concentration of reactant known---- can determine reactant concentration at any time
Initial and final reactant concentrations known---can determine rate constant

39. 2nd Order Integrated Rate Law
Rate = -Δ[A] = k [A]2 Δt
-take equation and integrate with calculus to get….
1 __ __ = kt
[A]t [A]0
[A]0 = initial concentration (t = 0)
[A]t = concentration after a period of time

40. Example 3: 2NO2(g)  2NO(g) + O2(g)
Using the data provided, find the rate constant if the rate law = k[NO2]2.
Time (s)
[NO2]
0.0
0.070
1.0 x 102
0.0150
2.0 x 102
0.0082
3.0 x 102
0.0057
k = 0.52 M-1s-1

41. Example 4:
NO2 reacts to form NO and O2 by second-order kinetics with a rate constant = 32.6 L/molmin. What is the [NO2] after 1 minute if the initial [NO2] = 0.15M?
[NO2] = 0.025M

42. Reaction Mechanism
Pathway or series of steps through which reactants converted to products
Not all steps proceed at the same rate
Sum of the steps = overall reaction
Rate of any step is directly proportional to the reactant concentrations in the step.

43. Ex. 1: 2NO(g) + O2(g)2NO2(g) Step 1: 2NO  N2O2
Step 2: N2O2 + O2  2NO2
__________________________
2NO + O2  2NO2

44. Rate-limiting Step
Step in a reaction mechanism that “limits” how fast products are formed.
“limits” the rate of reactants converting to products

45. Intermediates
Chemical compounds FORMED and CONSUMED in a reaction mechanism
Appear on both sides of chemical equation
Transition compounds between reactants and products
Unstable, only exist a short time

46. Determining the rate law for a reaction using a reaction mechanism
Overall Reaction Rate Law
Only determine experimentally
Overall Reaction Rate Law can be applied to find the steps in a reaction mechanism
Rate law indicates slowest step in the mechanism
Determine the fast steps remaining in the mechanism

47. Rate Laws for Elementary Steps in Mechanism
Predictable
Can use coefficients ONLY for these steps
aA + bB  dD + eE Rate = k[A]a [B]b
Overall reaction rate = rate of SLOWEST step

48. Ex. 2: NO2 + CO NO + CO2 Step 1: 2NO2  NO + NO3 (slow)
Step 2: NO3 + CO  NO2 + CO2 (fast)
Determine the rate law for each step.

49. Ex.3 Based on the following reaction mechanism….
F2 + NO2  NO2F + F (slow)
NO2 + F  NO2F (fast)
Write the overall reaction.
Determine the rate law for each step
What is the rate law for the overall reaction?

50. Catalysts
Chemical compounds (atoms, molecules, ions) that increase ONLY the reaction rate.
Increases both forward and reverse rates for a chemical reaction
Not consumed in the reaction, not altered
Present at the start and end of reaction
Have no effect on equilibrium constants (K), ΔH, or ΔS

51. Catalysts (cont.)
Change the PATH a chemical reaction takes to get reactants to products
Lowers activation energy (Ea) or energy needed for the reaction to start
Only small amount of the compound needed to increase the rate for a reaction with a lot of reactant
Recycle and Reuse

52. Catalyst Reaction Pathway

53. Catalysts (cont.) Example: H2O2 Decomposition Catalyst Ea (kJ/mol)
Reaction Rate
None
75.3 1
Catalase 8
6.3 x 1011