Chem 1B
Dr. White
Chapter 12: Chemical Kinetics
INTRODUCTION:
1. When considering a chemical reaction, one must ask if a reaction will go
(____________________) and how fast the reaction will go (_______________).
2. Kinetics Experiment:
3. Example: Oxidation of glucose
4. Factors affecting the rates of reaction
5. Why study rates of reactions?
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12.1 Reaction Rates
A. Rate of reaction:
1. Consider: NO2 (g) + CO (g) → NO (g) + CO2 (g)
Can measure the changes in concentration of NO:
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Example: For the reaction. 4 NO (g) + O2 (g) → 2 N2O3 (g), the [NO] is
decreasing at a rate of 1.60 x 10 -4 M/s, how fast is [O2] changing?
12.2 Rate Laws: An Introduction
A. Rate Law (Differential Rate Law) –
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4 12.3 Determining the Form of the Rate Law
A. Method of Initial Rates
Example: For the reaction below
2 HgCl2 (aq) + C2O4
data were collected:
1-
(aq) → 2 Cl
1-
(aq) + 2 CO
2
(g) + Hg2Cl2 (s) the following kinetic
Experiment [HgCl2] (M)
[C2O4 1-] (M)
1
0.105
0.15
2
0.105
0.30
3
0.052
0.30
What is the rate law? What is the value of k?
Initial Rate (M/min)
1.8 x 10-5
7.1 x 10-5
3.5 x 10-5
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5 Example: For the reaction:
H2O2 (aq) + 3I1- (aq) + 2H+ (aq) → I31- (aq) + 2 H2O (l) the following kinetic data were
collected:
Experiment [H2O2] (M)
[I1-] (M)
[H+]
1
0.010
0.010
0.00050
2
0.020
0.010
0.00050
3
0.010
0.020
0.00050
4
0.010
0.010
0.00100
What is the rate law and the value of the rate constant?
Initial Rate (M/s)
1.15 x 10-6
2.30 x 10-6
2.30 x 10-6
1.15 x 10-6
Example: Consider the following reaction: 2A + B → C + D. What are the units of k for
the reaction? Use the data below:
Trial
1
2
3
[A] (M)
0.225
0.320
0.225
[B] (M)
0.150
0.150
0.250
Rate (M/s)
0.0217
0.0439
0.0362
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12.4 Integrated Rate Laws -
A. First Order Integrated Rate Law
B. Second Order Integrated Rate Law
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C. Zero-Order Integrated Rate Law
D. Examples using integrated rate law
1. The rate law for a reaction is: rate = k[N2O5] and the rate constant is
4.80 x 10-4/s at 45°C.
a. If the initial concentration is 1.65 x 10-2 M, what is the
concentration after 825 s?
b. what fraction has decomposed in 825s?
c. How long would it take for the concentration of N2O5 to decrease
to 1.00 x 10-2 M?
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8 2. Determine the rate law for the reaction N2O5(g) →4 NO2(g) + O2(g) using
the data below.
3. The graphs below were obtained for the decomposition of NO2 at 298K
2 NO2 (g) → N2 (g) + O2 (g)
a. What is the rate law?
b. What is the value for the rate constant of this reaction (including
units)?
c. What is the initial concentration of NO2?
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9 -3
1000
0
ln [NO2]
600
200
400
600
800
400
200
y = -0.0029x - 3.91
R2 = 0.9999
-5
-6
-7
0
-200
0
200
400
1000
-4
y = 0.739x - 33.0
R2 = 0.8401
600
800
1000
-8
1200
time (seconds)
time (seconds)
0.025
y = -2.1E-05x + 0.017
R2 = 0.872
0.02
[NO2] (M)
1/[NO2] M-1
800
0.015
0.01
0.005
0
-0.005
0
200
400
600
800
time (seconds)
E. Half Life of a Reaction
1000
1200
1200
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Example: The following is a first order reaction:
SO2Cl2(g) → SO2 (g) + Cl2 (g)
At 320°C, the rate constant is 2.20x10-5/s.
a. What is the half life at this temp?
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b. How long (in hours) would it take for 50% of the sample to
decompose?
c. How long would it take for 75% to decompose?
d. How long would it take for 35% of the sample to decompose?
Kinetics and Chemical Equilibrium
Consider A+BC+D (all aq)
1. How do the rates of the forward and reverse rates compare at equilibrium?
2. K =
3. Suppose the forward and reverse rates are both second order overall. Write the
rate laws:
4. How does K relate to the rate laws?
5. Equilibrium Constants in kinetic terms:
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12 12.6 Reaction Mechanisms –
A. Elementary Reactions
Consider the following reaction:
NO2 (g) + CO (g) → NO (g) + CO2 (g)
Overall reaction
Proposed mechanism:
NO2 + NO2 → NO3 + NO (elementary reaction)
NO3 + CO → NO2 + CO2 (elementary reaction)
B. Proposed mechanisms must meet two requirements:
Even with these restrictions, generating a mechanism is difficult and relies on a lot of
guesswork. Our goal here is to be able to determine if a proposed mechanism is consistent
with experimental data, at least in simple cases.
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13 C. Rate Laws for Elementary Reactions
Example (mechanism with a slow initial step): For the reaction, O3 (g) + 2NO2
(g) → O2 (g) + N2O5 (g), the experimental rate law was determined to be: Rate
= k[O3][NO2]. Is the proposed mechanism below plausible?
proposed mechanism: O3 + NO2 → NO3 + O2 (slow)
NO3 + NO2 → N2O5 (fast)
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Dr. White
14 Example (mechanism with a fast initial step): The rate law for the following
reaction:
2 NO (g) + 2H2 (g) → N2 (g) + 2H2O (g)
was found to be: rate = k[NO]2[H2]. Is the following mechanism plausible for
this overall reaction?
i. 2NO (g)
N2O2 (g)
(fast, equilibrium)
ii. N2O2 (g) + H2 (g) → N2O (g) + H2O (g)
(slow)
iii. N2O (g) + H2 (g) → N2 (g) + H2O (g)
(fast)
Example: Consider the reaction: H2 (g) + I2 (g) → 2HI (g). The observed
rate law is first order with respect to each reactant. Show that the
proposed mechanism below fits the observed rate law.
i. I2 (g)
2 I (g)
ii. H2 (g) + 2I → 2HI (g)
(fast, equilibrium)
(slow)
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12.7 A Model for Chemical Kinetics
A. The Collision Model of Chemical Kinetics
B. Transition State Theory
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Example – reaction of CH3Br with OH-:
C. Molecular Orientation also Important
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D. Arrhenius Equation –
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18 Example: The relationship between rate constant and temperature for a second
order reaction is shown in the graph below:
-5
3.05E-03 3.10E-03 3.15E-03 3.20E-03 3.25E-03 3.30E-03 3.35E-03 3.40E-03
-5.5
-6
ln k
-6.5
-7
y = -1.07E+04x + 2.66E+01
-7.5
-8
-8.5
-9
-9.5
-10
1/T (1/K)
a. Calculate the activation energy for this reaction
b. Calculate the value of the rate constant at 500K.
Example: The rate constant for a first order reaction is 9.16x10 -3 s-1 at 0°C. Ea is
88.0 kJ/mol. If the temperature is raised 2°C, what is the new rate constant?
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Example: For the reaction C2H5I ---> C2H4 + HI
At 600K, k = 1.6 X 10-5 s-1. At 700K, k = 6.36 X10-3 s-1 What is Ea?
E. Relationship to Le Chatelier’s Principle
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20 12.8 Catalysis
A. Catalyst –
B. Heterogeneous Catalysts –
1. The Metal-Catalyzed Hydrogenation of Ethylene: H2C=CH2 + H2 → H3C−CH3
C. Homogeneous catalyst –
1. Example: production of sulfuric acid
uncatalyzed mechanism
S + O2 → SO2 (fast)
SO2 + 1/2 O2 → SO3 (slow)
SO3 + H2O → H2SO4 (fast)
catalyzed mechanism
S + O2 → SO2 (fast)
NO2 + SO2 → NO + SO3 (faster)
NO + 1/2 O2 → NO2 (faster)
SO3 + H2O → H2SO4 (fast)
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21 Example: Consider the following mechanism and indicate the species acting
as a catalyst.
i. Cl2  2 Cl
ii. N2O + Cl → N2 + ClO
iii. ClO + ClO → Cl2 + O2
D. Energy diagram of catalyzed and uncatalyzed reaction