15 February 2012
1

Objective: You will be able to:
 define “kinetics” and identify factors
that affect the rate of a reaction.
 write rate expressions for balanced
chemical reactions.
Agenda
Do now
II. Kinetics notes
III. Reaction Rates Demonstrations
IV. Rate constant and reaction rates
problems.
Homework: p. 602 #2, 3, 5, 7, 12, 13, 15,
16, 18: Thurs.
I.
3
Chemical Kinetics
Aspects of Chemistry
4
How can we predict whether or not a
reaction will take place?
 Thermodynamics
 Once started, how fast does the reaction
proceed?
 Chemical kinetics: this unit!
 How far will the reaction go before it
stops?
 Equilibrium: next unit

Chemical Kinetics


The area of chemistry concerned with the
speeds, or rates, at which a chemical reaction
occurs.
reaction rate: the change in the concentration
of a reactant or product with time (M/s)
 Why do reactions have such very different
rates?
 Steps in vision: 10-12 to 10-6 seconds!
 Graphite to diamonds: millions of years!
 In chemical industry, often more important to
maximize the speed of a reaction, not
necessarily yield.
A
D[A]
rate = Dt
D[B]
rate =
Dt
6
B
Chemical Kinetics
Reaction rate is the change in the concentration of a
reactant or a product with time (M/s).
A
B
D[A]
rate = Dt
D[A] = change in concentration of A over
time period Dt
D[B]
rate =
Dt
D[B] = change in concentration of B over
time period Dt
Because [A] decreases with time, D[A] is negative.
red-brown
2Br- (aq) + 2H+ (aq) + CO2 (g)
Br2 (aq) + HCOOH (aq)
time
t 1< t 2 < t 3
393 nm
light
8
Detector
D[Br2] a D Absorption
Br2 (aq) + HCOOH (aq)
2Br- (aq) + 2H+ (aq) + CO2 (g)
slope of
tangent
slope of
tangent
slope of
tangent
[Br2]final – [Br2]initial
D[Br2]
average rate = =Dt
tfinal - tinitial
9
instantaneous rate = rate for specific instance in time
Factors that Affect Reaction Rates
10




Concentration of reactants: higher
concentrations = faster reactions
 as concentration increases, the frequency of
collisions increases, increasing reaction rate
Temperature: increasing temperature increases
reaction rate because of increased KE
Physical state of reactants: homogeneous
mixtures of either liquids or gases react faster than
heterogeneous mixtures
Presence of a catalyst: affects the kinds of
collisions that lead to a reaction.
Question and Demo
11

Mine explosions from the ignition of
powdered coal dust are relatively
common, yet lumps of coal burn without
exploding. Explain.
Reaction Rates and Stoichiometry
2A
B
Two moles of A disappear for each mole of B that is formed.
1 D[A]
rate = 2 Dt
aA + bB
D[B]
rate =
Dt
cC + dD
1 D[A]
1 D[B]
1 D[C]
1 D[D]
rate = ==
=
a Dt
b Dt
c Dt
d Dt
12
Example
Write the rate expression for the following
reaction:
 CH4 (g) + 2O2 (g)
CO2 (g) + 2H2O (g)

Write the rate expression for the following reaction:
CH4 (g) + 2O2 (g)
CO2 (g) + 2H2O (g)
D[CH4]
D[CO2]
1 D[O2]
1 D[H2O]
rate = =
==
Dt
Dt
Dt
2 Dt
2
14
Practice Problems

Write the rate expressions for the following
reactions in terms of the disappearance of the
reactants and appearance of products.
a. I-(aq) + OCl-(aq)  Cl-(aq) + OI-(aq)
b. 4NH3(g) + 5O2(g)  4NO(g) + 6H2O(g)
rate a [Br2]
rate = k [Br2]
rate
= rate constant
k=
[Br2]
= 3.50 x 10-3 s-1
Using Rate Expressions
Consider the reaction:
 4NO2(g) + O2(g)  2N2O5(g)
Suppose that, at a particular moment during
the reaction, molecular oxygen is reacting
at the rate of 0.024 M/s.
a. At what rate is N2O5 being formed?
b. At what rate is NO2 reacting?
16 February 2012
Objective: You will be able to:
 solve rate expressions.
 determine the order of a reaction from
experimental data
Homework Quiz: N2(g) + 3H2(g) → 2NH3(g)
Suppose that at a particular moment during
the reaction, hydrogen is reacting at the rate of
0.074 M/s.
a. At what rate is NH3 being formed?
b. At what rate is nitrogen reacting?
18

Agenda
Do now
II. Iodine clock reaction.
III. Solving rate equations
IV. Determining order of reactions
Homework: p. 602 #15, 16, 18, 19, 20:
Mon after break
Hint: Use pressure just like concentration.
Diagnostic test (Tues after break)
I.
20
Example
Consider the reaction:
4PH3(g)  P4(g) + 6H2(g)
Suppose that, at a particular moment during
the reaction, molecular hydrogen is being
formed at the rate of 0.078 M/s.
a. At what rate is P4 being formed?
b. At what rate is PH3 reacting?
Problem
22
Consider the reaction between gaseous
hydrogen and gaseous nitrogen to produce
ammonia gas.
 At a particular time during the reaction,
H2(g) disappears at the rate of 3.0 M/s.
a. What is the rate of disappearance of
N2(g)?
b. What is the rate of appearance of NH3(g)?

23

If ammonia appears at 2.6 M/s, how fast
does hydrogen disappear?
The Rate Law
The rate law is a mathematical relationship that shows how
rate of reaction depends on the concentrations of reactants
aA + bB
cC + dD
Rate = k [A]x[B]y
x and y are small whole numbers that
relate to the number of molecules of A
and B that collide and are determined
experimentally!
The Rate Law
aA + bB
cC + dD
Rate = k [A]x[B]y
Reaction is xth order in A
Reaction is yth order in B
Reaction is (x +y)th order overall
Rate = k [A]1[B]2
Example
26
Experiment [A](M)
[B](M)
1
2
3
0.10
0.20
0.20

0.10
0.10
0.20
Rate = −d[A]/dt (M/s)
0.04
0.08
0.32
What is the numerical value of the rate
constant for the reaction described in the
table above? Specify units.
F2 (g) + 2ClO2 (g)







2FClO2 (g)
rate = k [F2]x[ClO2]y
Double [F2] with [ClO2] constant
Rate doubles
rate = k [F2][ClO2]
x=1
Quadruple [ClO2] with [F2] constant
Rate quadruples
y=1
Write the reaction rate expressions for the
following in terms of the disappearance of
the reactants and the appearance of
products:
a) 2H2(g) + O2(g)  2H2O(g)
b) 4NH3(g) + 5O2(g)  4NO(g) + 6H2O(g)
Consider the reaction
N2(g) + 3H2(g)  2NH3(g)
Suppose that at a particular moment during
the reaction molecular hydrogen is
reacting at a rate of 0.074 M/s.
a) At what rate is ammonia being formed?
b) At what rate is molecular nitrogen
reacting?
27 February 2012
Take Out: p. 602 #15, 16, 18, 19, 20
 Objective: You will be able to determine the
rate of a reaction given experimental data and
reactant concentrations.
 Homework Quiz: What is the rate law for
the reaction shown below?
 What is the rate when [A]=1.50 M and
[B]=0.50 M?

30
Run # Initial [A] ([A]0)
Initial [B] ([B]0)
Initial Rate (v0)
1
1.00 M
1.00 M
1.25 x 10-2 M/s
2
1.00 M
2.00 M
2.5 x 10-2 M/s
3
2.00 M
2.00 M
2.5 x 10-2 M/s
Agenda
31
Homework Quiz
II. Homework answers
III. Determining and solving rate laws
IV. Hand back tests and assignments
Homework: Diagnostic test
revisit/correct p. 603 #15, 16, 18
I.
Rate Laws
• Rate laws are always determined
experimentally.
• Reaction order is always defined in terms of
reactant (not product) concentrations.
• The order of a reactant is not related to the
stoichiometric coefficient of the reactant in the
balanced chemical equation.
F2 (g) + 2ClO2 (g)
2FClO2 (g)
rate = k [F2][ClO2] 1
32
Determine the rate law and calculate the rate constant for the
following reaction from the following data:
S2O82- (aq) + 3I- (aq)
2SO42- (aq) + I3- (aq)
Experiment
[S2O82-]
[I-]
Initial Rate
(M/s)
1
0.08
0.034
2.2 x 10-4
2
0.08
0.017
1.1 x 10-4
3
0.16
0.017
2.2 x 10-4
Determine the rate law and calculate the rate constant for the
following reaction from the following data:
S2O82- (aq) + 3I- (aq)
2SO42- (aq) + I3- (aq)
Experiment
[S2O82-]
[I-]
Initial Rate
(M/s)
1
0.08
0.034
2.2 x 10-4
2
0.08
0.017
1.1 x 10-4
3
0.16
0.017
2.2 x 10-4
rate = k [S2O82-]x[I-]y
y=1
x=1
rate = k [S2O82-][I-]
Double [I-], rate doubles (experiment 1 & 2)
Double [S2O82-], rate doubles (experiment 2 & 3)
34
2.2 x 10-4 M/s
rate
k=
=
= 0.08/M•s
2[S2O8 ][I ] (0.08 M)(0.034 M)
Practice Problems
The reaction of nitric oxide with hydrogen at
1280oC:
2NO(g) + 2H2(g)  N2(g) + 2H2O(g)
From the following data collected at this
temperature, determine (a) the rate law, (b)
the rate constant and (c) the rate of the
reaction when [NO] = 12.0x10-3 M and [H2]
= 6.0x10-3 M

35
Experiment
[NO] M
[H2] M
Initial Rate (M/s)
1
5.0x10-3
2.0x10-3
1.3x10-5
2
10.0x10-3
2.0x10-3
5.0x10-5
3
10.0x10-3
4.0x10-3
10.0x10-5
Calculate the rate of the reaction at the
time when [F2] = 0.010 M and [ClO2] =
0.020 M.
 F2(g) + 2ClO2(g)  2FClO2(g)

[F2] (M)
[ClO2] (M)
Initial Rate (M/s)
0.10
0.010
1.2x10-3
0.10
0.040
4.8x10-3
0.20
0.010
2.4x10-3
Consider the reaction X + Y  Z
From the following data, obtained at 360 K,
a) determine the order of the reaction
b) determine the initial rate of disappearance
of X when the concentration of X is 0.30 M
and that of Y is 0.40 M
37
Initial Rate of Disappearance
of X (M/s)
[X] (M)
[Y] (M)
0.053
0.10
0.50
0.127
0.20
0.30
1.02
0.40
0.60
0.254
0.20
0.60
0.509
0.40
0.30
38
Consider the reaction A B.
The rate of the reaction is 1.6x10-2 M/s
when the concentration of A is 0.35 M.
Calculate the rate constant if the reaction
is
a. first order in A
b. second order in A

The rate laws can be used to determine the
concentrations of any reactants at any
time during the course of a reaction.
29 Nov. 2010

40


Take Out Homework p. 603 #19, 21, 22, 23, 2529
Objective: SWBAT compare 1st order, 2nd order,
and zero order reactions, and describe how
temperature and activation energy effect the rate
constant.
Do now: Calculate the half life of the reaction
F2(g) + 2ClO2(g)  2FClO2(g), with rate data shown
[ClO2] (M)
Initial Rate (M/s)
below: [F2] (M)
0.10
0.010
1.2x10-3
0.10
0.040
4.8x10-3
0.20
0.010
2.4x10-3
28 February 2012
Take Out: Diagnostic test
 Objective: You will be able to determine
order of a reaction and k graphically.
 Homework Quiz: What is the rate law for
the reaction shown below?
 What is the rate when [A]=1.50 M and
[B]=0.50 M?

41
Run #
Initial [A] ([A]0)
Initial [B] ([B]0)
Initial Rate (v0)
1
1.00 M
1.00 M
1.25 x 10-2 M/s
2
1.00 M
2.00 M
2.5 x 10-2 M/s
3
2.00 M
2.00 M
2.5 x 10-2 M/s
Agenda
42
I.
II.
III.
Homework Quiz
1st order reactions graphically
Half life calculations
Homework: p. 603 #19, 20 (use Excel!),
24, 26
First Order (Overall) Reactions
43
rate depends on the concentration of a
single reactant raised to the first power.
rate = k[A] =  DA
Dt



Using calculus, this rate law is
transformed into an equation for a line:
ln[A] = ln[A]0 - kt
First-Order Reactions
A
k=
product
D[A]
rate = Dt
rate
M/s
=
= 1/s or s-1
M
[A]
[A] = [A]0e−kt
rate = k [A]
D[A]
= k [A]
Dt
ln[A] = ln[A]0 - kt
Graphical Determination of k
2N2O5
4NO2 (g) + O2 (g)
A non-graphical example
46

The reaction 2A
B is first order in A
with a rate constant of 2.8 x 10-2 s-1 at
800C. How long will it take for A to
decrease from 0.88 M to 0.14 M ?
The reaction 2A
B is first order in A with a rate constant
of 2.8 x 10-2 s-1 at 800C. How long will it take for A to decrease
from 0.88 M to 0.14 M ?
[A]0 = 0.88 M
ln[A] = ln[A]0 - kt
[A] = 0.14 M
kt = ln[A]0 – ln[A]
ln[A]0 – ln[A]
=
t=
k
47
ln
[A]0
[A]
k
ln
=
0.88 M
0.14 M
2.8 x
10-2
s-1
= 66 s
The conversion of cyclopropane to propene in the
gas phase is a first order reaction with a rate
constant of 6.7x10-4 s-1 at 500oC.
a)
b)
c)
If the initial concentration of cyclopropane was
0.25 M, what is the concentration after 8.8
minutes?
How long, in minutes, will it take for the
concentration of cyclopropane to decrease
from 0.25 M to 0.15 M?
How long, in minutes, will it take to convert
74% of the starting material?
29 February 2012
Objective: You will be able to:
 calculate the half-life of a first order reaction
 explore the relationship between time and
concentration of a second order reaction
Homework Quiz:
The conversion of cyclopropane to propene in the
gas phase is a first order reaction with a rate
constant of 6.7x10-4 s-1 at 500oC.
If the initial concentration of cyclopropane was
0.25 M, what is the concentration after 8.8
minutes?
49
The rate of decomposition of azomethane (C2H6N2)
is studied by monitoring partial pressure of the
reactant as a function of time:
CH3-N=N-CH3(g) → N2(g) + C2H6(g)
The data obtained at 300oC are shown here:
Time (s)
Partial Pressure of Azomethane (mmHg)
0
284
100
220
150
193
200
170
250
150
300
132
Are these values consistent with first-order kinetics? If
so, determine the rate constant.

51
The following gas-phase reaction was studied at
290oC by observing the change in pressure as a
function of time in a constant-volume vessel:
 ClCO2CCl3(g)  2COCl2(g)
 Determine the order of the reaction and the
rate constant based on the following data,
where P is the total pressure
Time (s)
P (mmHg)
0
400
2,000
316
4,000
248
6,000
196
8,000
155
10,000
122
Ethyl iodide (C2H5I) decomposes at a certain
temperature in the gas phase as follows:
C2H5I(g) → C2H4(g) + HI(g)
From the following data, determine the order of the
reaction and the rate constant:
Time (min)
[C2H5I] (M)
0
0.36
15
0.30
30
0.25
48
0.19
75
0.13
First-Order Reactions
The half-life, t½, is the time required for the concentration of a
reactant to decrease to half of its initial concentration.
t½ = t when [A] = [A]0/2
ln
t½ =
[A]0
[A]0/2
k
ln 2
0.693
=
=
k
k
What is the half-life of N2O5 if it decomposes with a rate constant
of 5.7 x 10-4 s-1?
How do you know decomposition is first order?
First-Order Reactions
The half-life, t½, is the time required for the concentration of a
reactant to decrease to half of its initial concentration.
t½ = t when [A] = [A]0/2
ln
t½ =
[A]0
[A]0/2
k
ln 2
0.693
=
=
k
k
What is the half-life of N2O5 if it decomposes with a rate constant
of 5.7 x 10-4 s-1?
0.693
t½ = ln 2 =
= 1200 s = 20 minutes
-4
-1
k
5.7 x 10 s
How do you know decomposition is first order?
54
units of k (s-1)
First-order reaction
A
product
# of
half-lives [A] = [A]0/n
1
2
55
2
4
3
8
4
16
56
The decomposition of ethane (C2H6) to
methyl radicals is a first-order reaction
with a rate constant of 5.36x10-4 s-1 at
700oC:
C2H6(g)  2CH3(g)
Calculate the half-life of the reaction in
minutes.

57

Calculate the half-life of the decomposition
of N2O5:
2N2O5  4NO2(g) + O2(g)
t (s)
[N2O5] (M)
ln [N2O5]
0
0.91
-0.094
300
0.75
-0.29
600
0.64
-0.45
1200
0.44
-0.82
3000
0.16
-1.83
Second-Order Reactions
A
product
D[A]
rate = Dt
rate
M/s
=
= 1/M•s
k=
2
2
M
[A]
1
1
=
+ kt
[A]
[A]0
58
D[A]
= k [A]2
Dt
[A] is the concentration of A at any time t
[A]0 is the concentration of A at time t=0
t½ = t when [A] = [A]0/2
1
t½ =
k[A]0
rate = k [A]2
Iodine atoms combine to form molecular
iodine in the gas phase:
I(g) + I(g)  I2(g)
This reaction follows second-order kinetics
and has the high rate constant 7.0x109/M·s
at 23oC.
a. If the initial concentration of I was 0.086
M, calculate the concentration after 2.0
minutes.
b. Calculate the half-life of the reaction if the
initial concentration of I is 0.60 M and if it
is 0.42 M.
The reaction 2A → B is second order with a rate
constant of 51/M·min at 24oC.
a. Starting with [A]o = 0.0092 M, how long
will it take for [A]t = 3.7x10-3 M?
b. Calculate the half-life of the reaction.
1 March 2012
61
Objective: You will be able to:
 determine the activation energy for a
reaction
 Homework Quiz:
The reaction 2A → B is second order with a rate
constant of 51/M·min at 24oC.
a. Starting with [A]o = 0.0092 M, how long will
it take for [A]t = 3.7x10-3 M?
b. Calculate the half-life of the reaction.

Agenda
62
Homework Quiz
II. Questions?
III. Kinetics Quiz
IV. Activation Energy
Homework: p.
I.
Zero-Order Reactions
A
product
D[A]
rate = Dt
D[A]
=k
Dt
rate
= M/s
k=
0
[A]
[A] = [A]0 - kt
[A] is the concentration of A at any time t
[A]0 is the concentration of A at time t = 0
t½ = t when [A] = [A]0/2
[A]0
t½ =
2k
63
rate = k [A]0 = k
Summary of the Kinetics of Zero-Order, First-Order
and Second-Order Reactions
Order
0
64
Rate Law
rate = k
1
rate = k [A]
2
[A]2
rate = k
Concentration-Time
Equation
[A] = [A]0 - kt
ln[A] = ln[A]0 - kt
1
1
=
+ kt
[A]
[A]0
Half-Life
t½ =
[A]0
2k
t½ = ln 2
k
1
t½ =
k[A]0
65
Activation Energy and Temperature
Dependence of Rate Constants
Reaction rates increase with increasing
temperature
 Ex: Hard boiling an egg
 Ex: Storing food
 How do reactions get started in the first
place?

Collision Theory
66

Chemical reactions occur as a result of
collisions between reacting molecules.
reaction rate depends on concentration
 But, the relationship is more complicated
than you might expect!
 Not all collisions result in reaction

Question
67

Explain in terms of collision theory why
temperature affects rate of reaction.
So, when does the reaction happen?
68
In order to react, colliding molecules must
have a total KE ≥ activation energy (Ea)
 Ea: minimum amount of energy required
to initiate a chemical reaction
 activated complex (transition state):
a temporary species formed by the
reactant molecules as a result of the
collision before they form the product.

A+B
Exothermic Reaction
+
AB+
C+D
Endothermic Reaction
The activation energy (Ea ) is the minimum amount of energy
required to initiate a chemical reaction.
=a barrier that prevents less energetic molecules from reacting
Rate Constant is Temp. Dependent
70
Arrhenius equation
k  Ae
( E a / RT )
Ea is the activation energy (J/mol)
R is the gas constant (8.314 J/K•mol)
T is the absolute temperature
A is the frequency factor
Alternate Arrhenius Equation
71

To relate k at two temperatures, T1 and T2:
The rate constants for the decomposition of
acetaldehyde:
CH3CHO(g) → CH4(g) + CO(g)
were measured at five different temperatures. The
data are shown below. Plot lnk versus 1/T, and
determine the activation energy (in kJ/mol) for the
3
reaction. (Note: the reaction
is 2 order in CH3CHO,
1
so k has the units of 1 / M 2  s )
1
2
k (1 / M  s)
0.011
0.035
T (K)
700
730
0.105
0.343
0.789
760
790
810
Determining
0.00
0.0012
0.00125
0.0013
0.00135
slope = -2.19x104
Ea
 slope = 
0.0014
0.00145
R
-1.00
-1.50
-2.00
ln K
Graphically

-0.50
-2.50
-3.00
-3.50
y = -21881x + 26.662
-4.00
-4.50
-5.00
Ea

R
1/T (K-1)
Determining activation energy
74
The second order rate constant for the
decomposition of nitrous oxide (N2O) into
nitrogen molecule and oxygen atom has been
measured at different temperatures. Determine
graphically the activation energy for the
reaction. k (1 / M  s)
T (oC)
1.87x10-3
0.0113
0.0569
0.244
600
650
700
750
5 March 2012
75
Objective: You will be able to:
 review and correct answers to the
multiple choice questions on the
diagnostic test.
 Homework Quiz:
 Please use the same sheet of paper as
last week!

Agenda
76
Homework Quiz
II. Homework answers
III. Correct and explain answers to
diagnostic test multiple choice questions.
Homework: Finish correcting and
explaining answers to multiple choice: due
Weds.
I.
With one partner:
77
Check your answers to the multiple choice
against my answers on the board.
 For each question you answered
incorrectly, or skipped, or guessed and
happened to get it right:
 Write 1 to 2 sentences to explain why the
correct answer is correct.
 Use resources! Textbook, notes,
internet…

7 March 2012
78
Objective: You will be able to:
 review, correct and explain answers to
the free response questions on the
diagnostic test.
 Do now: Look at your free response 1-6
and decide on your first three preferences
for creating a poster and explaining your
answers. Write them down on your slip of
paper.

Agenda
79
I.
II.
Objective and agenda
Correct and explain answers to
diagnostic test free response questions
With your group…
80
1.
2.
3.
4.
Check your answers with the answer key.
Make notes about how to solve the
problem/answer the question.
Design and create a poster that shows the work
and answers, as well as additional
explanations of how to solve the problem or
answer the question.
Post your poster in the room! Then, go look at
other groups posters and correct your work.
30 Nov. 2010
Take Out Homework p. 605# 31, 32, 35,
37, 39
 Objective: SWBAT use the Arrhenius
equation to solve for rate constants and
temperatures, and solve practice problems
on kinetics.
 Do now: Match

81
Order Rate Law
2
rate = k[A]
Conc-Time Eq. Half Life Eq.
[A]=[A]0-kt
t1/2=1/k[A]o
1
rate = k[A]2 1/[A]=1/[A]0 + kt
0
rate = k
ln[A]=ln[A]0 –kt
t1/2=ln2/k
t1/2=[A]0/2k
Agenda
82
Homework solutions
II. Using the Arrhenius equation part 2
III. Molecular orientation
IV. Problem Set work time
Homework: Complete problem set and
p. 605 #40, 42
Quiz tomorrow
I.
8 March 2012
83
Objective: You will be able to:
 review, correct and explain answers to
the free response questions on the
diagnostic test.
 describe the reaction mechanism of a
reaction
 Do now: Finish and hang up your poster.
(10 min.)

Agenda
84
Objective and agenda
II. Gallery Walk: Correct and explain
answers to diagnostic test free response
questions
III. Using the Alternate Arrhenius Equation
IV. Hand back quizzes
Homework p. 605 #44, 45, 49, 51, 52, 54:
Mon.
I.
Gallery Walk
85
Walk with your group
 Spend 5 minutes at each station
 Correct/complete your work and make
notes of how/why each problem is solved.

Using the alternate Arrhenius Equation
86

The rate constant of a first order reaction
is 3.46x10-2 /s at 298 K. What is the rate
constant at 350 K if the activation energy
for the reaction is 50.2 kJ/mol?
Using the Arrhenius Equation
87

The first order rate constant for the
reaction of methyl chloride (CH3Cl) with
water to produce methanol (CH3OH) and
hydrochloric acid (HCl) is 3.32x10-10/s at
25oC. Calculate the rate constant at 40oC
if the activation energy is 116 kJ/mol.
88
Frequency of Collisions and
Orientation Factor
For simple reactions (between atoms, for
example) the frequency factor (A) is
proportional to the frequency of collision
between the reacting species.
 “Orientation factor” is also important.

Importance of Molecular Orientation
effective collision
ineffective collision
89
Reaction Mechanisms
90
A balanced chemical equation doesn’t tell
us much about how the reaction actually
takes place.
 It may represent the sum of elementary
steps
 Reaction mechanism: the sequence of
elementary steps that leads to product
formation.

Reaction Mechanisms
The overall progress of a chemical reaction can be represented
at the molecular level by a series of simple elementary steps
or elementary reactions.
The sequence of elementary steps that leads to product
formation is the reaction mechanism.
2NO (g) + O2 (g)
2NO2 (g)
N2O2 is detected during the reaction!
91
Elementary step:
NO + NO
N 2O 2
+ Elementary step:
N2O2 + O2
2NO2
Overall reaction:
2NO + O2
2NO2
2NO (g) + O2 (g)
Mechanism:
92
2NO2 (g)
13 March 2012
93

Objective: You will be able to
 identify overall reactions, intermediates
and rate laws for reaction mechanisms.
Agenda
94
Objectives and Agenda
II. Review: Reaction mechanisms
III. Elementary step examples
IV. Catalysts
Homework: p. 605 #44, 45, 49, 51, 52, 54,
55, 56, 61: Tues.
I.
Intermediates are species that appear in a reaction
mechanism but not in the overall balanced equation.
An intermediate is always formed in an early elementary step
and consumed in a later elementary step.
Elementary step:
NO + NO
N 2O 2
+ Elementary step:
N2O2 + O2
2NO2
Overall reaction:
2NO + O2
2NO2
The molecularity of a reaction is the number of molecules
reacting in an elementary step.
•
Unimolecular reaction – elementary step with 1 molecule
•
Bimolecular reaction – elementary step with 2 molecules
•
Termolecular reaction – elementary step with 3 molecules
95
Rate Laws and Elementary Steps
Unimolecular reaction
A
products
rate = k [A]
Bimolecular reaction
A+B
products
rate = k [A][B]
Bimolecular reaction
A+A
products
rate = k [A]2
Writing plausible reaction mechanisms:
•
The sum of the elementary steps must give the overall
balanced equation for the reaction.
•
The rate-determining step should predict the same rate
law that is determined experimentally.
The rate-determining step is the slowest step in the
sequence of steps leading to product formation.
96
The experimental rate law for the reaction between NO2 and CO
to produce NO and CO2 is rate = k[NO2]2. The reaction is
believed to occur via two steps:
Step 1:
NO2 + NO2
NO + NO3
Step 2:
NO3 + CO
NO2 + CO2
What is the equation for the overall reaction?
What is the intermediate?
What can you say about the relative rates of steps 1 and 2?
The experimental rate law for the reaction between NO2 and CO
to produce NO and CO2 is rate = k[NO2]2. The reaction is
believed to occur via two steps:
Step 1:
NO2 + NO2
NO + NO3
Step 2:
NO3 + CO
NO2 + CO2
What is the equation for the overall reaction?
NO2+ CO
NO + CO2
What is the intermediate?
NO3
What can you say about the relative rates of steps 1 and 2?
98
rate = k[NO2]2 is the rate law for step 1 so
step 1 must be slower than step 2
Rate Determining Step
99

rate determining step: the slowest step in
the sequence of steps leading to product
formation.
Problem
100

Propose a mechanism for the overall
reaction:
2A + 2B → A2B2
Example
The gas-phase decomposition of nitrous oxide
(N2O) is believed to occur via two elementary
steps:
Step 1:
N2O  N2 + O
Step 2 N2O + O  N2 + O2
Experimentally the rate law is found to be
rate = k[N2O].
a) Write the equation for the overall reaction.
b) Identify the intermediates.
c) What can you say about the relative rates of
steps 1 and 2?

101
102
NO2 + F2 → NO2F + F
NO2 + F → NO2F
a.
b.
c.
d.
Write the overall reaction.
What is the intermediate?
If the rate law is rate = k[NO2][F2], which
step is the rate determining step?
Which step proceeds at the fastest rate?
103
Hydrogen and iodine monochloride react
as follows:
H2(g) + 2ICl(g) → 2HCl(g) + I2(g)
The rate law for the reaction is
rate = k[H2][ICl]. Suggest a possible
mechanism for the reaction.

Decomposition of Hydrogen Peroxide
104
2H2O2(aq)  2H2O(l) + O2(g)
Can be catalyzed using iodide ions (I-)
rate = k[H2O2][I-] Why?!
Determined experimentally.
Step 1: H2O2 + I-  H2O + IOStep 2: H2O2 + IO-  H2O + O2 + I-
105
For the decomposition for H2O2, the
reaction rate depends on the
concentration of I- ions, even though Idoesn’t appear in the overall equation.
 I- is a catalyst for the reaction.

A catalyst is a substance that increases the rate of a
chemical reaction without itself being consumed.
k  A  e ( Ea / RT )
Ea
Catalyzed
k
Uncatalyzed
ratecatalyzed > rateuncatalyzed
106
Ea′ < Ea
Catalysts
107
forms an alternative reaction pathway
 lowers overall activation energy
 for example, it might form an
intermediate with the reactant.
 Ex: 2KClO3(s)  2KCl(s) + 3O2(g)
Very slow, until you add MnO2, a catalyst.
The MnO2 can be recovered at the end of
the reaction!

Week of March 12
108
Step 1:
HBr + O2 → HOOBr
Step 2: HOOBr + HBr → 2HOBr
Step 3:
HOBr + HBr → H2O + Br2
Step 4:
HOBr + HBr → H2O + Br2
a. Write the equation for the overall reaction.
b. Identify the intermediate(s).
c. What can you say about the relative rate of
each step if the rate law is rate =
k[HBr][O2]?
13 March 2012
109
Objective: You will be able to
 identify and describe the effect of
catalysts in a reaction mechanism.
 Agenda:
I.
Homework Quiz
II. Homework Answers
III. Catalysts
IV. Problem Set
Homework: Problem Set: Monday

Catalyst Example: Ozone Cycle
110
Step 1: O2(g) + hv → O(g) + O(g)
 Step 2: O(g) + O2(g) → O3(g)
 Step 3: O3(g) + hv → O2(g) + O(g)
 Step 4: O(g) + O(g) → O2(g)
 Overall: O3(g) + O2(g) → O2(g) + O3(g)
This cycle continually repeats, producing
and destroying ozone at the same rate
while absorbing harmful ultraviolet light
from the sun.

Chlorofluorocarbons and Ozone
Chlorine atoms from CFCs released into the
atmosphere catalyze the O3(g) → O2(g)
reaction.
 Net result: ozone is depleted faster that is
generated by the natural cycle.
 Cl atoms from CFCs deplete the ozone layer!
 Step 1: 2Cl(g) + 2O3(g) → 2ClO(g) + 2O2(g)
 Step 2: ClO(g) + ClO(g) → O2(g) + 2Cl(g)
 Overall: 2O3(g) → 3O2(g)

111
In heterogeneous catalysis, the reactants and the catalysts
are in different phases (usually, catalyst is a solid, reactants
are gases or liquids).
•
Haber synthesis of ammonia
•
Ostwald process for the production of nitric acid
•
Catalytic converters
In homogeneous catalysis, the reactants and the catalysts
are dispersed in a single phase, usually liquid.
112
•
Acid catalysis
•
Base catalysis
Haber Process
Synthesis of Ammonia
Extremely slow at room temperature. Must be fast
and high yield!
Process occurs on the surface of the Fe/Al2O3/K2O
catalyst, which weakens the covalent N-N and H-H
bonds.
Fe/Al2O3/K2O
N2 (g) + 3H2 (g)
catalyst
113
2NH3 (g)
Ostwald Process
4NH3 (g) + 5O2 (g)
Pt catalyst
2NO (g) + O2 (g)
2NO2 (g) + H2O (l)
114
4NO (g) + 6H2O (g)
2NO2 (g)
HNO2 (aq) + HNO3 (aq)
Pt-Rh catalysts used
in Ostwald process
Catalytic Converters
catalytic
CO + Unburned Hydrocarbons + O2 converter
CO2 + H2O
catalytic
2NO + 2NO2 converter
2N2 + 3O2
115
Enzyme Catalysis
biological catalysts
116
Binding of Glucose to Hexokinase
117
14 March 2012
118
Objective: You will be able to:
 demonstrate your knowledge of
chemical kinetics on a problem set and a
lab.
 Agenda:
I.
Objectives and Agenda
II. Work time:
I. Problem Set
II. Kinetics Pre-Lab

AP Exam
119
Monday, May 7
 If you have a year average >80%, you pay
$13 (full cost = $87!)
 This is due, in CASH (no coins), by next
Friday.
 If your average is <80%, I’ll chat with you
privately today about your options.

Homework
120
Pre-lab: due tomorrow
 Lab procedure: read by tomorrow
 Problem set: due Monday
 Kinetics test: Tuesday

Expectations
121
Choose ONE person to work with.
 Work either on the problem set or the prelab questions (or split your time…)
 Stay at your table.
 Use a professional tone and volume of
voice.
 Use this time wisely!

15 March 2012
122
Sit at a lab table with your group.
 Take Out: Lab notebook and lab packet
 Objective: You will be able to:
 determine the rate law and the
activation energy for the oxidation of
iodide ions by bromate ions in the
presence of an acid.

Homework
123
Problem Set due Monday
 Kinetics Unit Test Tuesday
 Gas Unit revisions due tomorrow

Logistics
124
Half of the groups will do Part 1 on page 5
while the other half does steps 1-3 on page
6.
 Then, we’ll switch!

Changes to the Procedure
125
Instead of reaction strips, you’ll be using
spot plates.
 Instead of inverting one reaction strip
over the other and shaking down to mix,
you’ll be adding the drops of KBrO3,
starting the stopwatch, and stirring with a
toothpick to mix.
 You must do this at the same way, in the
same order, and at the same speed each
time!

126
Put the reagents for reaction strip 1 in one
well plate.
 If more than 2 drops of KBrO3, place the
drops in a second well plate.
 Transfer them with a separate pipette so
you can dispense them all at once into
the first well plate.
 Start timing and stir.

Precision and Consistency
127
Be very precise in your work, or your
results won’t be meaningful.
 Be very consistent in the way your carry
out the procedure: the way you hold the
pipette to drop solutions, the way you add
the KBrO3 (from “reaction strip 2”), the
rate at which you stir, when you start and
stop timing, etc.

Reagents and Equipment
128

Leave reagents at the front table. Bring
your labeled pipettes to the table to fill
them.
Data
129

Record your data immediately and
carefully in tables in your lab notebook.
19 March 2012
130


Objective: You will be able to:
 determine the reaction order, rate law,
and activation energy for an iodine clock
reaction.
Reminder: $13 (cash) due by Friday for
AP Exam
Homework
131
Problem Set due today
 Kinetics Test tomorrow
 10 MC
 1-2 FRQ

What’s the purpose?
132
22 March 2012
133

Objective: You will be able to:
 determine the rate law, reaction
constant and activation energy for the
iodine clock reaction.
Agenda
134
Finish lab
II. Clean up/return materials
III. Work on lab calculations, analysis and
conclusions in your lab notebook
 Note: all data, etc. must also be in your lab
notebook!
Homework: Lab notebook due Monday
$13 for AP Exam due by 8:00 am TOMORROW!!!
I.
Water baths
135
Warm water bath (40oC) on the side bench.
 If it’s too cool, remove some water, and add
some hot water from the beaker on the hot
plate.
 It should be shallow! Don’t swamp your
spot plate. Record the actual temp.
 Ice bath (OoC): create one using ice and
water in your metal pan. Use a little
thermometer to record the temperature.

Safety
136
Keep your goggles on your eyes!
 One warning
 Then you’re out.
 Label your reagents and store them
carefully.
 Use a professional voice and stay at your
table unless you need to get something.

Cleanup
137
Keep your labeled pipettes in the cassette
case.
 Rinse transfer pipettes in water and squirt
out water to dry.
 Return equipment to the cart.
 Make sure your lab table is clean and neat.
