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15 February 2012

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Title: 15 February 2012


1
15 February 2012
  • 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.

2
Agenda
  • Do now
  • Kinetics notes
  • Reaction Rates Demonstrations
  • Rate constant and reaction rates problems.
  • Homework p. 602 2, 3, 5, 7, 12, 13, 15, 16, 18
    Thurs.

3
Chemical Kinetics
4
Aspects of Chemistry
  • 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

5
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.

6
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7
Chemical Kinetics
Reaction rate is the change in the concentration
of a reactant or a product with time (M/s).
DA change in concentration of A over
time period Dt
DB change in concentration of B over
time period Dt
Because A decreases with time, DA is
negative.
8
red-brown
t1lt t2 lt t3
DBr2 a D Absorption
9
instantaneous rate rate for specific instance
in time
10
Factors that Affect Reaction Rates
  • 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.

11
Question and Demo
  • Mine explosions from the ignition of powdered
    coal dust are relatively common, yet lumps of
    coal burn without exploding. Explain.

12
Reaction Rates and Stoichiometry
Two moles of A disappear for each mole of B that
is formed.
13
Example
  • Write the rate expression for the following
    reaction
  • CH4 (g) 2O2 (g) CO2 (g) 2H2O (g)

14
Write the rate expression for the following
reaction CH4 (g) 2O2 (g) CO2 (g)
2H2O (g)
15
Practice Problems
  • Write the rate expressions for the following
    reactions in terms of the disappearance of the
    reactants and appearance of products.
  • I-(aq) OCl-(aq) ? Cl-(aq) OI-(aq)
  • 4NH3(g) 5O2(g) ? 4NO(g) 6H2O(g)

16
rate a Br2
rate k Br2
rate constant
3.50 x 10-3 s-1
17
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.
  • At what rate is N2O5 being formed?
  • At what rate is NO2 reacting?

18
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.
  • At what rate is NH3 being formed?
  • At what rate is nitrogen reacting?

19
Agenda
  • Do now
  • Iodine clock reaction.
  • Solving rate equations
  • 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)

20







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21
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.
  • At what rate is P4 being formed?
  • At what rate is PH3 reacting?

22
Problem
  • 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.
  • What is the rate of disappearance of N2(g)?
  • What is the rate of appearance of NH3(g)?

23
  • If ammonia appears at 2.6 M/s, how fast does
    hydrogen disappear?

24
The Rate Law
The rate law is a mathematical relationship that
shows how rate of reaction depends on the
concentrations of reactants
Rate k AxBy
x and y are small whole numbers that relate to
the number of molecules of A and B that collide
and are determined experimentally!
25
The Rate Law
Rate k AxBy
Reaction is xth order in A
Reaction is yth order in B
Reaction is (x y)th order overall
Rate k A1B2
26
Example
Experiment A(M) B(M) Rate -dA/dt (M/s)
1 0.10 0.10 0.04
2 0.10 0.20 0.08
3 0.20 0.20 0.32
  • What is the numerical value of the rate constant
    for the reaction described in the table above?
    Specify units.

27
  • rate k F2xClO2y
  • Double F2 with ClO2 constant
  • Rate doubles
  • x 1
  • Quadruple ClO2 with F2 constant
  • Rate quadruples
  • y 1

rate k F2ClO2
28
  • Write the reaction rate expressions for the
    following in terms of the disappearance of the
    reactants and the appearance of products
  • 2H2(g) O2(g) ? 2H2O(g)
  • 4NH3(g) 5O2(g) ? 4NO(g) 6H2O(g)

29
  • 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.
  • At what rate is ammonia being formed?
  • At what rate is molecular nitrogen reacting?

30
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 A1.50 M and B0.50 M?

Run Initial A (A0) Initial B (B0) 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
31
Agenda
  • Homework Quiz
  • Homework answers
  • Determining and solving rate laws
  • Hand back tests and assignments
  • Homework Diagnostic test
  • revisit/correct p. 603 15, 16, 18

32
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.

rate k F2ClO2
33
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
34
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-xI-y
y 1
x 1
rate k S2O82-I-
Double I-, rate doubles (experiment 1 2)
Double S2O82-, rate doubles (experiment 2 3)
0.08/Ms
35
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

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
36
  • 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
37
  • Consider the reaction X Y ? Z
  • From the following data, obtained at 360 K,
  • determine the order of the reaction
  • 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

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
  • first order in A
  • second order in A

39
  • The rate laws can be used to determine the
    concentrations of any reactants at any time
    during the course of a reaction.

40
29 Nov. 2010
  • Take Out Homework p. 603 19, 21, 22, 23, 25-29
  • 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 below

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
41
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 A1.50 M and B0.50 M?

Run Initial A (A0) Initial B (B0) 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
42
Agenda
  • Homework Quiz
  • 1st order reactions graphically
  • Half life calculations
  • Homework p. 603 19, 20 (use Excel!), 24, 26

43
First Order (Overall) Reactions
  • rate depends on the concentration of a single
    reactant raised to the first power.
  • rate kA
  • Using calculus, this rate law is transformed into
    an equation for a line

lnA lnA0 - kt
44
First-Order Reactions
rate k A
k
1/s or s-1
45
Graphical Determination of k
46
A non-graphical example
  • 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 ?

47
A0 0.88 M
lnA lnA0 - kt
A 0.14 M
kt lnA0 lnA
t
66 s
48
  • 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?
  • 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?

49
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?

50
  • The rate of decomposition of azomethane (C2H6N2)
    is studied by monitoring partial pressure of the
    reactant as a function of time
  • CH3-NN-CH3(g) ? N2(g) C2H6(g)
  • The data obtained at 300oC are shown here
  • Are these values consistent with first-order
    kinetics? If so, determine the rate constant.

Time (s) Partial Pressure of Azomethane (mmHg)
0 284
100 220
150 193
200 170
250 150
300 132
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
52
  • 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
53
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 A0/2
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?
54
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 A0/2
What is the half-life of N2O5 if it decomposes
with a rate constant of 5.7 x 10-4 s-1?
1200 s 20 minutes
How do you know decomposition is first order?
units of k (s-1)
55
First-order reaction
1
2
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
58
Second-Order Reactions
rate k A2
k
1/Ms
A is the concentration of A at any time t
A0 is the concentration of A at time t0
t½ t when A A0/2
59
  • 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/Ms at 23oC.
  • If the initial concentration of I was 0.086 M,
    calculate the concentration after 2.0 minutes.
  • Calculate the half-life of the reaction if the
    initial concentration of I is 0.60 M and if it is
    0.42 M.

60
  • The reaction 2A ? B is second order with a rate
    constant of 51/Mmin at 24oC.
  • Starting with Ao 0.0092 M, how long will it
    take for At 3.7x10-3 M?
  • Calculate the half-life of the reaction.

61
1 March 2012
  • 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/Mmin at 24oC.
  • Starting with Ao 0.0092 M, how long will it
    take for At 3.7x10-3 M?
  • Calculate the half-life of the reaction.

62
Agenda
  • Homework Quiz
  • Questions?
  • Kinetics Quiz
  • Activation Energy
  • Homework p.

63
Zero-Order Reactions
rate k A0 k
k
M/s
A is the concentration of A at any time t
A A0 - kt
A0 is the concentration of A at time t 0
t½ t when A A0/2
64
Summary of the Kinetics of Zero-Order,
First-Order and Second-Order Reactions
A A0 - kt
0
rate k
1
rate k A
lnA lnA0 - kt
2
rate k A2
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?

66
Collision Theory
  • 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

67
Question
  • Explain in terms of collision theory why
    temperature affects rate of reaction.

68
So, when does the reaction happen?
  • 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.

69
Exothermic Reaction
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
70
Rate Constant is Temp. Dependent
Arrhenius equation
Ea is the activation energy (J/mol)
R is the gas constant (8.314 J/Kmol)
T is the absolute temperature
A is the frequency factor
71
Alternate Arrhenius Equation
  • To relate k at two temperatures, T1 and T2

72
  • 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 reaction. (Note the reaction is
    order in CH3CHO, so k has the units of
    )

k T (K)
0.011 700
0.035 730
0.105 760
0.343 790
0.789 810
73
Determining Graphically
  • slope -2.19x104
  • slope

74
Determining activation energy
  • 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 T (oC)
1.87x10-3 600
0.0113 650
0.0569 700
0.244 750
75
5 March 2012
  • 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!

76
Agenda
  • Homework Quiz
  • Homework answers
  • Correct and explain answers to diagnostic test
    multiple choice questions.
  • Homework Finish correcting and explaining
    answers to multiple choice due Weds.

77
With one partner
  • 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

78
7 March 2012
  • 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.

79
Agenda
  1. Objective and agenda
  2. Correct and explain answers to diagnostic test
    free response questions

80
With your group
  1. Check your answers with the answer key.
  2. Make notes about how to solve the problem/answer
    the question.
  3. 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.
  4. Post your poster in the room! Then, go look at
    other groups posters and correct your work.

81
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

Order Rate Law Conc-Time Eq. Half Life Eq.
2 rate kA AA0-kt t1/21/kAo
1 rate kA2 1/A1/A0 kt t1/2ln2/k
0 rate k lnAlnA0 kt t1/2A0/2k
82
Agenda
  • Homework solutions
  • Using the Arrhenius equation part 2
  • Molecular orientation
  • Problem Set work time
  • Homework Complete problem set and
  • p. 605 40, 42
  • Quiz tomorrow

83
8 March 2012
  • 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.)

84
Agenda
  • Objective and agenda
  • Gallery Walk Correct and explain answers to
    diagnostic test free response questions
  • Using the Alternate Arrhenius Equation
  • Hand back quizzes
  • Homework p. 605 44, 45, 49, 51, 52, 54 Mon.

85
Gallery Walk
  • Walk with your group
  • Spend 5 minutes at each station
  • Correct/complete your work and make notes of
    how/why each problem is solved.

86
Using the alternate Arrhenius Equation
  • 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?

87
Using the Arrhenius Equation
  • 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.

89
Importance of Molecular Orientation
effective collision
ineffective collision
90
Reaction Mechanisms
  • A balanced chemical equation doesnt 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.

91
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.
N2O2 is detected during the reaction!
92
Mechanism
93
13 March 2012
  • Objective You will be able to
  • identify overall reactions, intermediates and
    rate laws for reaction mechanisms.

94
Agenda
  • Objectives and Agenda
  • Review Reaction mechanisms
  • Elementary step examples
  • Catalysts
  • Homework p. 605 44, 45, 49, 51, 52, 54, 55, 56,
    61 Tues.

95
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.
  • 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

96
Rate Laws and Elementary Steps
rate k A
rate k AB
rate k A2
  • 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.
97
The experimental rate law for the reaction
between NO2 and CO to produce NO and CO2 is rate
kNO22. The reaction is believed to occur via
two steps
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?
98
The experimental rate law for the reaction
between NO2 and CO to produce NO and CO2 is rate
kNO22. The reaction is believed to occur via
two steps
What is the equation for the overall reaction?
What is the intermediate?
NO3
What can you say about the relative rates of
steps 1 and 2?
rate kNO22 is the rate law for step 1 so
step 1 must be slower than step 2
99
Rate Determining Step
  • rate determining step the slowest step in the
    sequence of steps leading to product formation.

100
Problem
  • Propose a mechanism for the overall reaction
  • 2A 2B ? A2B2

101
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 kN2O.
  • Write the equation for the overall reaction.
  • Identify the intermediates.
  • What can you say about the relative rates of
    steps 1 and 2?

102
  • NO2 F2 ? NO2F F
  • NO2 F ? NO2F
  • Write the overall reaction.
  • What is the intermediate?
  • If the rate law is rate kNO2F2, 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 kH2ICl. Suggest a possible mechanism
    for the reaction.

104
Decomposition of Hydrogen Peroxide
  • 2H2O2(aq) ? 2H2O(l) O2(g)
  • Can be catalyzed using iodide ions (I-)
  • rate kH2O2I- Why?!
  • Determined experimentally.
  • Step 1 H2O2 I- ? H2O IO-
  • Step 2 H2O2 IO- ? H2O O2 I-

105
  • For the decomposition for H2O2, the reaction rate
    depends on the concentration of I- ions, even
    though I- doesnt appear in the overall equation.
  • I- is a catalyst for the reaction.

106
A catalyst is a substance that increases the rate
of a chemical reaction without itself being
consumed.
Uncatalyzed
Catalyzed
ratecatalyzed gt rateuncatalyzed
107
Catalysts
  • 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!

108
Week of March 12
  • Step 1 HBr O2 ? HOOBr
  • Step 2 HOOBr HBr ? 2HOBr
  • Step 3 HOBr HBr ? H2O Br2
  • Step 4 HOBr HBr ? H2O Br2
  • Write the equation for the overall reaction.
  • Identify the intermediate(s).
  • What can you say about the relative rate of each
    step if the rate law is rate kHBrO2?

109
13 March 2012
  • Objective You will be able to
  • identify and describe the effect of catalysts in
    a reaction mechanism.
  • Agenda
  • Homework Quiz
  • Homework Answers
  • Catalysts
  • Problem Set
  • Homework Problem Set Monday

110
Catalyst Example Ozone Cycle
  • 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.
  • hv ultraviolet light

111
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)

112
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.
  • Acid catalysis
  • Base catalysis

113
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.
114
Ostwald Process
Pt catalyst
115
Catalytic Converters
116
Enzyme Catalysis biological catalysts
117
Binding of Glucose to Hexokinase
118
14 March 2012
  • Objective You will be able to
  • demonstrate your knowledge of chemical kinetics
    on a problem set and a lab.
  • Agenda
  • Objectives and Agenda
  • Work time
  • Problem Set
  • Kinetics Pre-Lab

119
AP Exam
  • Monday, May 7
  • If you have a year average gt80, you pay 13
    (full cost 87!)
  • This is due, in CASH (no coins), by next Friday.
  • If your average is lt80, Ill chat with you
    privately today about your options.

120
Homework
  • Pre-lab due tomorrow
  • Lab procedure read by tomorrow
  • Problem set due Monday
  • Kinetics test Tuesday

121
Expectations
  • Choose ONE person to work with.
  • Work either on the problem set or the pre-lab
    questions (or split your time)
  • Stay at your table.
  • Use a professional tone and volume of voice.
  • Use this time wisely!

122
15 March 2012
  • 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.

123
Homework
  • Problem Set due Monday
  • Kinetics Unit Test Tuesday
  • Gas Unit revisions due tomorrow

124
Logistics
  • Half of the groups will do Part 1 on page 5 while
    the other half does steps 1-3 on page 6.
  • Then, well switch!

125
Changes to the Procedure
  • Instead of reaction strips, youll be using spot
    plates.
  • Instead of inverting one reaction strip over the
    other and shaking down to mix, youll 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.

127
Precision and Consistency
  • Be very precise in your work, or your results
    wont 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.

128
Reagents and Equipment
  • Leave reagents at the front table. Bring your
    labeled pipettes to the table to fill them.

129
Data
  • Record your data immediately and carefully in
    tables in your lab notebook.

130
19 March 2012
  • 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

131
Homework
  • Problem Set due today
  • Kinetics Test tomorrow
  • 10 MC
  • 1-2 FRQ

132
Whats the purpose?
133
22 March 2012
  • Objective You will be able to
  • determine the rate law, reaction constant and
    activation energy for the iodine clock reaction.

134
Agenda
  • Finish lab
  • Clean up/return materials
  • 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 800 am TOMORROW!!!

135
Water baths
  • Warm water bath (40oC) on the side bench.
  • If its too cool, remove some water, and add some
    hot water from the beaker on the hot plate.
  • It should be shallow! Dont 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.

136
Safety
  • Keep your goggles on your eyes!
  • One warning
  • Then youre out.
  • Label your reagents and store them carefully.
  • Use a professional voice and stay at your table
    unless you need to get something.

137
Cleanup
  • 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.
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