15 February 2012

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.

1

Agenda

I. Do nowII. Kinetics notesIII. Reaction Rates DemonstrationsIV. Rate constant and reaction rates

problems.Homework: p. 602 #2, 3, 5, 7, 12,

13, 15, 16, 18: Thurs.

Chemical Kinetics3

Aspects of Chemistry4

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.

6

A B

rate = -[A]t

rate = [B]t

Chemical KineticsReaction rate is the change in the concentration of a reactant or a product with time (M/s).

A B

rate = -[A]t

rate = [B]t

[A] = change in concentration of A over time period t

[B] = change in concentration of B over time period t

Because [A] decreases with time, [A] is negative.

8

Br2 (aq) + HCOOH (aq) 2Br- (aq) + 2H+ (aq) + CO2 (g)

time

393 nmlight

Detector

[Br2] Absorption

red-brown

t1< t2 < t3

9

Br2 (aq) + HCOOH (aq) 2Br- (aq) + 2H+ (aq) + CO2 (g)

average rate = -[Br2]t

= -[Br2]final – [Br2]initial

tfinal - tinitial

slope oftangent

slope oftangent slope of

tangent

instantaneous rate = rate for specific instance in time

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.

10

Question and Demo

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

11

12

Reaction Rates and Stoichiometry2A B

Two moles of A disappear for each mole of B that is formed.

rate = [B]t

rate = -[A]t

12

aA + bB cC + dD

rate = -[A]t

1a

= -[B]t

1b

=[C]t

1c

=[D]t

1d

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)

rate = -[CH4]

t= -

[O2]t

12

=[H2O]

t12

=[CO2]

t

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 [Br2]

rate = k [Br2]

k = rate[Br2]

= rate constant

= 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 dataHomework 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?

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Agenda

I. Do nowII. Iodine clock reaction.III. Solving rate equationsIV. Determining order of reactionsHomework: p. 602 #15, 16, 18, 19, 20:

Mon after breakHint: Use pressure just like concentration.Diagnostic test (Tues after break)

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

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)?

22

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

23

The Rate LawThe 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 LawaA + bB cC + dD

Rate = k [A]x[B]y

Reaction is xth order in AReaction is yth order in BReaction is (x +y)th order overall

Rate = k [A]1[B]2

Example

Experiment

[A](M) [B](M) Rate = −d[A]/dt (M/s)

1 0.10 0.10 0.042 0.10 0.20 0.083 0.20 0.20 0.32

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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][ClO2]

rate = k [F2]x[ClO2]y

Double [F2] with [ClO2] constant Rate doubles 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 reactionN2(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/s2 1.00 M 2.00 M 2.5 x 10-2 M/s3 2.00 M 2.00 M 2.5 x 10-2 M/s

Agenda

I. Homework QuizII. Homework answersIII. Determining and solving rate lawsIV. Hand back tests and assignmentsHomework: Diagnostic testrevisit/correct p. 603 #15, 16, 18

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32

F2 (g) + 2ClO2 (g) 2FClO2 (g)

rate = k [F2][ClO2]

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.

1

Determine the rate law and calculate the rate constant for the following reaction from the following data:S2O8

2- (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

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Determine the rate law and calculate the rate constant for the following reaction from the following data:S2O8

2- (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

Double [I-], rate doubles (experiment 1 & 2)

y = 1

Double [S2O82-], rate doubles (experiment 2 & 3)

x = 1

k = rate

[S2O82-][I-]

=2.2 x 10-4 M/s

(0.08 M)(0.034 M)= 0.08/M•s

rate = k [S2O82-][I-]

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

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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 ZFrom the following data, obtained at 360 K, a) determine the order of the reactionb) 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

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Initial Rate of Disappearance of X (M/s)

[X] (M) [Y] (M)

0.053 0.10 0.500.127 0.20 0.301.02 0.40 0.600.254 0.20 0.600.509 0.40 0.30

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 isa. first order in Ab. second order in A

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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 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:

40

[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

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/s3 2.00 M 2.00 M 2.5 x 10-2 M/s

Agenda

I. Homework QuizII. 1st order reactions graphicallyIII. Half life calculations

Homework: p. 603 #19, 20 (use Excel!), 24, 26

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First Order (Overall) Reactions

rate depends on the concentration of a single reactant raised to the first power.

rate = k[A] =

Using calculus, this rate law is transformed into an equation for a line:

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tA

ln[A] = ln[A]0 - kt

First-Order Reactions

A product rate = -[A]t

rate = k [A]

k = rate[A]

= 1/s or s-1M/sM=

[A]t = k [A]-

[A] = [A]0e−ktln[A] = ln[A]0 - kt

2N2O5 4NO2 (g) + O2 (g)

Graphical Determination of k

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 ?

46

47

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 ?

ln[A] = ln[A]0 - kt

kt = ln[A]0 – ln[A]

t =ln[A]0 – ln[A]

k= 66 s

[A]0 = 0.88 M

[A] = 0.14 M

ln[A]0

[A]k

=ln

0.88 M0.14 M

2.8 x 10-2 s-1=

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) If the initial concentration of cyclopropane was 0.25 M, what is the concentration after 8.8 minutes?

b) How long, in minutes, will it take for the concentration of cyclopropane to decrease from 0.25 M to 0.15 M?

c) 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:

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

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

51

Time (s) P (mmHg)0 4002,000 3164,000 2486,000 1968,000 15510,000 122

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[A]0

[A]0/2k

=t½ln 2k

=0.693

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?

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] = [A]0/2

ln[A]0

[A]0/2k

=t½ln 2k

=0.693

k=

What is the half-life of N2O5 if it decomposes with a rate constant of 5.7 x 10-4 s-1?

t½ln 2k

=0.693

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

A product

First-order reaction

# of half-lives [A] = [A]0/n

1

2

3

4

2

4

8

16

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.

56

Calculate the half-life of the decomposition of N2O5:

2N2O5 4NO2(g) + O2(g)

57

t (s) [N2O5] (M) ln [N2O5]0 0.91 -0.094300 0.75 -0.29600 0.64 -0.451200 0.44 -0.823000 0.16 -1.83

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Second-Order Reactions

A product rate = -[A]t

rate = k [A]2

k = rate[A]2 = 1/M•sM/s

M2=[A]t = k [A]2-

[A] is the concentration of A at any time t[A]0 is the concentration of A at time t=0

1[A]

=1

[A]0+ kt

t½ = t when [A] = [A]0/2

t½ = 1k[A]0

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

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.

61

Agenda

I. Homework QuizII. Questions?III. Kinetics QuizIV. Activation EnergyHomework: p.

62

63

Zero-Order Reactions

A product rate = -[A]t

rate = k [A]0 = k

k = rate[A]0 = M/s

[A]t = k-

[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

t½ = [A]0

2k

[A] = [A]0 - kt

64

Summary of the Kinetics of Zero-Order, First-Orderand Second-Order Reactions

Order Rate LawConcentration-Time

Equation Half-Life

0

1

2

rate = k

rate = k [A]

rate = k [A]2

ln[A] = ln[A]0 - kt

1[A]

=1

[A]0+ kt

[A] = [A]0 - kt

t½ln 2k

=

t½ = [A]0

2k

t½ = 1k[A]0

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?

65

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

66

Question

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

67

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.

68

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

A + B AB C + D++

Rate Constant is Temp. Dependent70

T is the absolute temperatureA is the frequency factor

Arrhenius equation

)/( RTEaeAk Ea is the activation energy (J/mol)

R is the gas constant (8.314 J/K•mol)

Alternate Arrhenius Equation

To relate k at two temperatures, T1 and T2:

71

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 )

23

sM 21

/1

k T (K)0.011 7000.035 7300.105 7600.343 7900.789 810

)/1( 21

sM

Determining Graphically

slope = -2.19x104

slope = REa

REa

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.

74

k T (oC)1.87x10-3 6000.0113 6500.0569 7000.244 750

)/1( sM

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!

75

Agenda

I. Homework QuizII. Homework answersIII. Correct and explain answers to

diagnostic test multiple choice questions.

Homework: Finish correcting and explaining answers to multiple choice: due Weds.

76

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…

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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.

78

Agenda

I. Objective and agendaII. Correct and explain answers to

diagnostic test free response questions

79

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.

80

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 Conc-Time Eq. Half Life Eq.2 rate = k[A] [A]=[A]0-kt t1/2=1/k[A]o

1 rate = k[A]2 1/[A]=1/[A]0 + kt t1/2=ln2/k

0 rate = k ln[A]=ln[A]0 –kt t1/2=[A]0/2k

Agenda

I. Homework solutionsII. Using the Arrhenius equation part 2III. Molecular orientationIV. Problem Set work timeHomework: Complete problem set and

p. 605 #40, 42Quiz tomorrow

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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.)

83

Agenda

I. Objective and agendaII. Gallery Walk: Correct and explain

answers to diagnostic test free response questions

III. Using the Alternate Arrhenius Equation

IV. Hand back quizzesHomework p. 605 #44, 45, 49, 51, 52,

54: Mon.

84

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.

85

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?

86

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.

87

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.

88

89

Importance of Molecular Orientation

effective collision

ineffective collision

Reaction Mechanisms

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.

90

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.

2NO (g) + O2 (g) 2NO2 (g)

N2O2 is detected during the reaction!

Elementary step: NO + NO N2O2

Elementary step: N2O2 + O2 2NO2

Overall reaction: 2NO + O2 2NO2

+

92

2NO (g) + O2 (g) 2NO2 (g)Mechanism:

13 March 2012

Objective: You will be able to identify overall reactions,

intermediates and rate laws for reaction mechanisms.

93

Agenda

I. Objectives and AgendaII. Review: Reaction mechanismsIII. Elementary step examplesIV. CatalystsHomework: p. 605 #44, 45, 49, 51,

52, 54, 55, 56, 61: Tues.

94

95

Elementary step: NO + NO N2O2

Elementary step: N2O2 + O2 2NO2

Overall reaction: 2NO + O2 2NO2

+

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

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

Rate Laws and Elementary Steps

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.

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?

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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?

rate = k[NO2]2 is the rate law for step 1 so step 1 must be slower than step 2

Rate Determining Step

rate determining step: the slowest step in the sequence of steps leading to product formation.

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Problem

Propose a mechanism for the overall reaction:

2A + 2B → A2B2

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Example The gas-phase decomposition of nitrous

oxide (N2O) is believed to occur via two elementary steps:Step 1: N2O N2 + OStep 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?

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NO2 + F2 → NO2F + FNO2 + F → NO2F

a. Write the overall reaction.b. What is the intermediate?c. If the rate law is rate = k[NO2][F2], which

step is the rate determining step? d. Which step proceeds at the fastest rate?

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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.

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Decomposition of Hydrogen Peroxide

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 + IO-

Step 2: H2O2 + IO- H2O + O2 + I-

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For the decomposition for H2O2, the reaction rate depends on the concentration of I- ions, even though I- doesn’t appear in the overall equation.

I- is a catalyst for the reaction.

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A catalyst is a substance that increases the rate of a chemical reaction without itself being consumed.

Ea k

ratecatalyzed > rateuncatalyzed

Ea < Ea′

UncatalyzedCatalyzed

)/( RTEaeAk

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!

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Week of March 12

Step 1: HBr + O2 → HOOBrStep 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]?

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13 March 2012

Objective: You will be able to identify and describe the effect of

catalysts in a reaction mechanism. Agenda:I. Homework QuizII. Homework AnswersIII.CatalystsIV.Problem SetHomework: Problem Set: Monday

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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

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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)

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In heterogeneous catalysis, the reactants and the catalysts are in different phases (usually, catalyst is a solid, reactants are gases or liquids).

In homogeneous catalysis, the reactants and the catalysts are dispersed in a single phase, usually liquid.

• Haber synthesis of ammonia

• Ostwald process for the production of nitric acid

• Catalytic converters

• Acid catalysis

• Base catalysis

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N2 (g) + 3H2 (g) 2NH3 (g)Fe/Al2O3/K2O

catalyst

Haber ProcessSynthesis 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.

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Ostwald Process

Pt-Rh catalysts usedin Ostwald process

4NH3 (g) + 5O2 (g) 4NO (g) + 6H2O (g)Pt catalyst

2NO (g) + O2 (g) 2NO2 (g)

2NO2 (g) + H2O (l) HNO2 (aq) + HNO3 (aq)

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Catalytic Converters

CO + Unburned Hydrocarbons + O2 CO2 + H2Ocatalytic

converter

2NO + 2NO2 2N2 + 3O2

catalyticconverter

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Enzyme Catalysisbiological catalysts

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Binding of Glucose to Hexokinase

14 March 2012

Objective: You will be able to: demonstrate your knowledge of

chemical kinetics on a problem set and a lab.

Agenda:I. Objectives and AgendaII.Work time:

I. Problem SetII.Kinetics Pre-Lab

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AP Exam

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.

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Homework

Pre-lab: due tomorrow Lab procedure: read by tomorrow Problem set: due Monday Kinetics test: Tuesday

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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!

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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.

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Homework

Problem Set due Monday Kinetics Unit Test Tuesday Gas Unit revisions due tomorrow

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Logistics

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!

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Changes to the Procedure

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!

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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.

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Precision and Consistency

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.

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Reagents and Equipment

Leave reagents at the front table. Bring your labeled pipettes to the table to fill them.

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Data

Record your data immediately and carefully in tables in your lab notebook.

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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

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Homework

Problem Set due today Kinetics Test tomorrow

10 MC 1-2 FRQ

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What’s the purpose?132

22 March 2012

Objective: You will be able to: determine the rate law, reaction

constant and activation energy for the iodine clock reaction.

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Agenda

I. Finish labII. Clean up/return materialsIII. 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!!!

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Water baths

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.

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Safety

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.

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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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