Chemical Kinetics Chapter 11

1
Chemical KineticsChapter 11
H2O2 decomposition in an insect
H2O2 decomposition catalyzed by MnO2
2
Chemical Kinetics

• KINETICS the study of REACTION RATES and their
  relation to the way the reaction proceeds, i.e.,
  its MECHANISM(steps in the process).
• We can use thermodynamics to tell if a reaction
  is product- or reactant-favored.
• Only KINETICS will tell us HOW FAST the reaction
  happens!

3
Meaning of Rate of Reaction

• rate the change per interval of time
• Exspeed (distance/ time) is a RATE

4
Reaction Rates

• Types of rates -
• initial rate
• average rate
• instantaneous rate

5
Determining a Reaction Rate

• Blue dye is oxidized with bleach.
• Its concentration decreases with time.
• The rate the change in dye conc with time can
  be determined from the plot.

NOTE THE UNITS OF RATE IN CHEMICAL REACTIONS!!!
6
Units

• Molarity moles/L
• Time hr, min, s
• Overall M/time or moles/L.time

7
C. Measurement
General reaction rate calculated by dividing
rate expressions by stoichiometric coefficients
8

• Example Calculate the rate of N2O5
  decomposition at 1.0 minute using this graph.
  How are the rates of formation of NO2 and O2
  related to the rate of decomposition of N2O5?

9
Determining a Reaction Rate
Instantaneous Rate is the slope of the tangent
line at a given point!
10
Three Types of Rates are NOT equal!
11
Example 2 Sample Calculation of Average Rate

• Rate of decomposition of N2O5(g) at 67oC
• Time(min) 0 1 2 3 4
• N2O5 0.160 0.113 0.080 0.056
  0.040
• What is the average rate of decomposition of
  dinitrogen pentoxide during the first three
  minutes?

12
II. Reaction Rate and Concentration

• A. Reactant concentration and collision theory
• Reactions occur when
• molecules collide to exchange or rearrange atoms
• Effective collisions occur when molecules have
  correct energy and orientation

13
Factors Affecting Rate

• Concentrations
• and physical state of reactants and
  products
• 2. Temperature
• 3. Catalysts
• Rate Law relates the rate of the reaction to the
  concentration of the reactants

14
Concentrations Rates
0.3 M HCl
6 M HCl
Mg(s) 2 HCl(aq) ---gt MgCl2(aq)
H2(g) As concentration increases, what happens to
rate?
15
Factors Affecting Rates

• Physical state of reactants

Greater surface area means greater rate of
reaction!
16
Factors Affecting Rates

• Temperature

Bleach at 54 C
Bleach at 22 C
As Temperature increases, what happens to rate?
Why?
17
Factors Affecting Rates

• Catalysts catalyzed decomp of H2O2
• Catalysts are substances that speed up a reaction
  but are unchanged by the reaction
• 2 H2O2 --gt 2 H2O O2

18
Effect of Concentration on Reaction Rate

• To propose a reaction mechanism, we study
• reaction rate and
• its concentration dependence

19
B. Rate Expression and Rate Constant
The rate law is R kAmBn
20
B. continued

• Rate k AmBnCp
• The exponents m, n, and p
• are the reaction order
• can be 0, 1, 2 or fractions
• must be determined by experiment!
• Overall Order Sum of m, n and p

21
Rate Constant, k

• Relates rate and concentration at a given
  temperature

Order Units of K
0 M Time-1
1 Time-1
2 M-1 Time-1
3 M-2 Time-1
General Formula M(1- overall order) time-1
22
-For one reactant

• aA? products
• Rate k Am
• m order of reaction with respect to A
• If m 0, zero order
• m 1, 1st order
• m 2, 2nd order

23

• If m 0, rxn. is zero order.
• Rate k A0
• If A doubles, rate DOES NOT CHANGE.
• If m 1, rxn. is 1st order in A
• Rate k A1
• If A doubles, then rate doubles
• If m 2, rxn. is 2nd order in A
• Rate k A2
• If A doubles, then rate quadruples

24
C. Determining Rates

• Use the equation
• Rate 2 A2m A2 m
• Rate 1 A1m A1

Pick any two points from the given data!
4.56 .195 m 3.15 .162
25

• The initial rate of decomposition of
  acetaldehyde, CH3CHO, was measured at a series
  of different concentrations and at a constant
  temperature.
• Using the data below, determine the order of the
  reaction (m) in the equation
• Rate kCH3CHOm
• CH3CHO(g) ? CH4(g) CO(g)

CH3CHO (mol/L) 0.162 0.195 0.273 0.410 0.518
Rate (mol/Lmin) 3.15 4.56 8.94 20.2 35.2
26
Deriving Rate Laws

• Rate of rxn k CH3CHO2
• Here the rate goes up by
• when initial conc. doubles.
• Therefore, we say this reaction is
  order.

FOUR
SECOND
27

• Consider the rate data for the decomposition
  of CH3CHO given in the above example. Knowing
  that the reaction is 2nd order, determine the
• a. the value of the rate constant (k)

28

• b. the rate of the reaction when CH3CHO
  0.452mol/L

29
Order of a Reaction for more than one reactant

• aA bB? products
• Rate kAmBn
• Overall order of reaction m n
• Rate 2 A2m B2 n
• Rate 1 A1mB1n

30
Sample Problem

• The data below are for the reaction of nitrogen
  (II) oxide with hydrogen at 800oC.
  2NO(g) 2H2(g) ? N2(g) 2H2O(g)
• Determine the order of the reaction with respect
  to both reactants, calculate the value of the
  rate constant, and determine the rate of
  formation of product when NO0.0024M and
  H20.0042M.
• Strategy Choose two experiments where conc. of
  one reactant is CONSTANT and other is changed
  solve for m and n separately!

31
Sample Problem

• Ex The initial rate of a reaction A B ? C was
  measured with the results below. State the rate
  law, the value of the rate constant, and the rate
  of reaction when
• A 0.050M and B 0.100M.

32
Sample Problem

• Ex The following data were collected for this
  reaction at constant temperature
• 2NO(g) Br2(g) ? 2NOBr(g)
• State the rate law, and determine the rate of
  this reaction when NO0.15M and Br20.25M.

33
Concentration/Time Relations

• What is concentration of reactant as function of
  time?
• Consider FIRST ORDER REACTIONS
• The rate law is

34
Concentration/Time Relations

• Integrating - (? A / ? time) k A, we get

A / A0 fraction remaining after time t has
elapsed.
Called the integrated first-order rate law.
35
Using the Integrated Rate Law

• The integrated rate law suggests a way to tell
  the order based on experiment.
• 2 N2O5(g) ---gt 4 NO2(g) O2(g)
• Time (min) N2O50 (M) ln N2O50
• 0 1.00 0
• 1.0 0.705 -0.35
• 2.0 0.497 -0.70
• 5.0 0.173 -1.75

36
Using the Integrated Rate Law
Plot of ln N2O5 vs. time is a straight line!
Data of conc. vs. time plot do not fit straight
line.
37
Using the Integrated Rate Law
Plot of ln N2O5 vs. time is a straight line!
Eqn. for straight line y mx b

• All 1st order reactions have straight line plot
  for ln A vs. time.
• (2nd order gives straight line for plot of 1/A
  vs. time)

38
Properties of Reactions
39
Half-Life

• HALF-LIFE is the time it takes for 1/2 a sample
  is disappear.
• For 1st order reactions, the concept of HALF-LIFE
  is especially useful.

Active Figure 15.9
40
Half-Life

• Reaction is 1st order decomposition of H2O2.

41
Half-Life

• Reaction after 1 half-life.
• 1/2 of the reactant has been consumed and 1/2
  remains.

42
Half-Life

• After 2 half-lives 1/4 of the reactant remains.

43
Half-Life

• A 3 half-lives 1/8 of the reactant remains.

44
Half-Life

• After 4 half-lives 1/16 of the reactant remains.

45
Half-Lives of Radioactive Elements

• Rate of decay of radioactive isotopes given in
  terms of 1/2-life.
• 238U --gt 234Th He 4.5 x 109 y
• 14C --gt 14N beta 5730 y
• 131I --gt 131Xe beta 8.05 d
• Element 106 - seaborgium263Sg 0.9 s

46
Activation Energy

• A. definition the minimum amount of energy
  required to react, Ea.
• activated complex - a short-lived molecule formed
  when reactants collide it can return to
  reactants or form products.

Reaction coordinate diagram
47
Effective Collisions

• How can the number of effective collisions be
  increased?
• Increase concentration of reactants,
• Increase temperature of reaction,
• Increase surface area of reactants.

48
Potential Energy Diagrams
Energy of Activated Complex
49
CATALYSIS

• Catalysis LOWERS activation energy!

MnO2 catalyzes decomposition of H2O2 2 H2O2 ---gt
2 H2O O2
50
CATALYSIS

• Used in auto exhaust systems Pt, NiO

2 CO O2 ---gt 2 CO2 2 NO ---gt N2 O2
51
CATALYSIS

• 2. Polymers H2CCH2 ---gt polyethylene
• 3. Acetic acid
• CH3OH CO --gt CH3CO2H
• 4. Enzymes biological catalysts

52
REACTION MECHANISMS

• Definition of mechanism how reactants are
  converted to products at the molecular level.
• Most chemical reactions DO NOT occur in a single
  step!

53
Reaction Mechanisms

• Most reactions involve a sequence of elementary
  steps (a single step in a reaction)
• The slow step in a reaction is called the rate
  determining step.
• Adding elementary
• steps
• gives NET or
• OVERALL reaction.

54
Reaction Mechanisms
Rate k I- H2O2 Step 1 slow HOOH I-
--gt HOI OH- Step 2 fast HOI I- --gt
I2 OH- Step 3 fast 2 OH- 2 H --gt 2
H2O Net 2 I- H2O2 2 H ---gt I2 2
H2O

• The species HOI and OH- are intermediates
  (produced in one elementary step but reacted in
  another).
• Intermediates do NOT appear in the net equation
  OR the rate law!

55
Reaction Mechanisms

• Br from biomass burning destroys stratospheric
  ozone.
• (See R.J. Cicerone, Science, volume 263, page
  1243, 1994.)
• Step 1 Br O3 ---gt BrO O2
• Step 2 Cl O3 ---gt ClO
  O2
• Step 3 BrO ClO light ---gt Br Cl O2
• NET 2 O3 ---gt 3 O2
• Br is a catalyst (is a reactant in an elementary
  step, but is unchanged at end of reaction)
• Catalysts do NOT appear in the net equation OR
  the rate law!

56
Sample Problem

• Cl2(g) ? 2Cl(g) Fast
• Cl(g) CHCl3(g) ? CCl3(g) HCl(g) Slow
• CCl3(g) Cl(g) ? CCl4(g) Fast

• Identify the following
• Rate determining step,
• Overall reaction
• If there are any intermediates or catalysts in
  the reaction, and what they are.

57
Sample Problem

• H2O2(aq) I1-(aq) ? H2O(l)
  IO1-(aq) Slow
• H2O2(aq) IO1-(aq) ? H2O(l) O2(g) I1
  (aq) Fast

• Identify the following
• Rate determining step,
• Overall reaction
• If there are any intermediates or catalysts in
  the reaction, and what they are.

58
Sample Problem

• O3(g) Cl(g) ? O2(g) ClO(g) Slow
• ClO(g) O(g) ? Cl(g) O2(g) Fast

• Identify the following
• Rate determining step,
• Overall reaction
• If there are any intermediates or catalysts in
  the reaction, and what they are.

59
Ozone Decomposition over Antarctica
2 O3 (g) ---gt 3 O2 (g)