Chapter 11 Chemical Kinetics:
Rates and Mechanisms of Chemical Reactions
Objectives
1.
2.
3.
4.
5.
Students will understand the meaning of the phase “rate of reaction” and will know how to measure the
rate of reaction in several different ways.
Students will be able to write the rate law for a chemical reaction from initial rates data.
Students will be able to distinguish between zero, first, second, and third order reactions.
Students will know a little of the theory of reaction kinetics based on collision theory.
Students will be introduced to reaction mechanisms, intermediates, and catalysis via energy diagrams
illustrating reaction rates.
I. Meaning of Reaction Rate
A. Definition
Kinetics-
•
•
We can use thermodynamics to tell if a reaction is product-or
reactant-favored.
Only KINETICS will tell us HOW FAST the reaction happens!
Rate-
Types of rates-
B. Units
C. Measurement
aA + bB  cC + dD
1
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?
Decomposition of N2O5
0.09
0.08
Molar Concentration of
N2O5
0.07
0.06
(mol/L)
0.05
0.04
0.03
0.02
0.01
0
0
20
40
60
80
100
Time (sec)
Example 2. Rate of decomposition of N2O5(g) at 67oC:
Time(min) 0
1
2
3
4
[N2O5]
0.160 0.113
0.080
0.056
.040
What is the average rate of decomposition of dinitrogen pentoxide during the first
three minutes?
II.
Reaction Rate and Concentration
A. Reactant concentration and collision theory


2
Factors affecting rate:
1.
2.
3.

Rate Law:
B. Rate Expression and Rate Constant
-
for one reactant
C. Determining Rates
3
Example: The initial rate of decomposition of acetaldehyde, CH3CHO, was
measured at a series of different concentrations and a constant temperature.
Using the data below, determine the order of the reaction (m) in the equation:
rate = k[CH3CHO]m
CH3CHO (g)  CH4 (g)
+ CO (g)
CH3CHO (mol/L)
0.162
0.195
0.273
0.410
0.518
Rate (mol/L*min)
3.15
4.56
8.94
20.2
35.2
Example: Consider the rate data for the decomposition of CH3CHO given in the
above example. Knowing that the reaction is 2nd order, determine the
a) rate constant (k)
b) rate of reaction when the concentration of CH3CHO is 0.452 mol/L.
4
D. Order of a Reaction for more than one reactant
Example: The data below are for the reaction of nitrogen (II) oxide with hydrogen at
800oC. Determine the order of reaction with respect to both H2 and NO. Determine the
rate constant and the initial rate when [NO] = 0.0024 M and [H2] = 0.0042 M.
2 NO (g) + 2 H2 (g)  N2 (g) + 2 H2O (g)
Experiment
1
2
3
4
5
6
Initial Conc. Initial Conc. Initial Rate of Formation
[NO] (mol/L) [H2] (mol/L)
of N2 (mol/L*min)
0.0010
0.0020
0.0030
0.0040
0.0040
0.0040
0.0040
0.0040
0.0040
0.0010
0.0020
0.0030
0.12
0.48
1.08
0.48
0.96
1.44
5
Example: The initial rate of reaction, A + B  C, was measured with the
results below. State the rate law, the value of the rate constant, and the rate of
the reaction when [A] = 0.050 M and [B] = 0.100 M.
Experiment [A] (mol/L) [B] (mol/L) Initial Rate (mol/L*s)
1
2
3
0.1
0.1
0.2
0.1
0.2
0.1
4.0 x 10-5
4.0 x 10-5
16.0 x 10-5
Example: The following data were collected for the formation of NOBr(g).
Determine the rate law, the value of the rate constant, and the rate of the
reaction when [NO] = 0.15 M and [Br2] = 0.25 M.
2 NO (g) + Br2 (g)  2 NOBr (g)
Experiment [NO] (mol/L) [Br2] (mol/L) Initial Rate (mol/L*s)
1
2
3
4
0.1
0.1
0.2
0.3
0.1
0.2
0.1
0.1
12
24
48
108
6
III. Reactant Concentration and Time
IV. Activation Energy
A. Definition

Activated Complex
B. Effective Collisions
C. Activation Energy Diagrams and Catalysts
7
A catalyst provides an alternate reaction pathway, which has lower activation energy
than an uncatalyzed reaction.
8
V. Reaction Mechanisms
A. Definition
B. Rate Determining Step
C. Intermediates
D. Catalysts
9
E. Examples:

Step 1:
H2O2(aq) + I-(aq)
Step 2:
H2O2(aq) + IO-(aq)
Step 1:
O3 (g) + Cl (g)

Step 2:
ClO(g) + O(g)

H2O +
 H2O +
IO-(aq)
O2 (aq)
O2 (g) +
Cl (g)
+ O2(g)
Slow
+ I-(aq)Fast
ClO(g)
Slow
Fast
10