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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. Agenda. Do now Kinetics notes Reaction Rates Demonstrations Rate constant and reaction rates problems.

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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.
agenda
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.

aspects of chemistry
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
chemical kinetics1
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.
slide6
A B

rate =

D[A]

D[B]

rate = -

Dt

Dt

slide7
A B

rate =

D[A]

D[B]

rate = -

Dt

Dt

Chemical Kinetics

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

D[A] = change in concentration of A over

time period Dt

D[B] = change in concentration of B over

time period Dt

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

slide8
Br2(aq) + HCOOH (aq) 2Br-(aq) + 2H+(aq) + CO2(g)

time

393 nm

Detector

light

red-brown

t1< t2 < t3

D[Br2] aD Absorption

slide9
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

instantaneous rate = rate for specific instance in time

factors that affect reaction rates
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.
question and demo
Question and Demo
  • Mine explosions from the ignition of powdered coal dust are relatively common, yet lumps of coal burn without exploding. Explain.
slide12
2A B

aA + bB cC + dD

rate = -

=

=

rate = -

= -

D[C]

D[B]

D[A]

D[B]

D[D]

D[A]

rate =

1

1

1

1

1

Dt

Dt

Dt

Dt

Dt

Dt

c

d

a

2

b

Reaction Rates and Stoichiometry

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

example
Example
  • Write the rate expression for the following reaction:
  • CH4 (g) + 2O2 (g) CO2 (g) + 2H2O (g)
slide14
D[CO2]

=

Dt

D[CH4]

rate = -

Dt

D[H2O]

=

Dt

D[O2]

= -

1

1

Dt

2

2

Write the rate expression for the following reaction:

CH4(g) + 2O2(g) CO2(g) + 2H2O (g)

practice problems
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)
slide16
rate

k =

[Br2]

rate a [Br2]

rate = k [Br2]

= rate constant

= 3.50 x 10-3 s-1

using rate expressions
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?
16 february 2012
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?
agenda1
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)

example1
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?
problem
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)?
slide24
aA + bB cC + dD

The Rate Law

The rate law is a mathematical relationship that shows how rate of reaction depends on the concentrations of reactants

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!

slide25
aA + bB cC + dD

The Rate Law

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

example2
Example
  • What is the numerical value of the rate constant for the reaction described in the table above? Specify units.
slide27
F2(g) + 2ClO2(g) 2FClO2(g)
  • rate = k [F2]x[ClO2]y
  • Double [F2] with [ClO2] constant
  • Rate doubles
  • x = 1
  • Quadruple [ClO2] with [F2] constant
  • Rate quadruples
  • y = 1

rate = k [F2][ClO2]

slide28
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)
slide29
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?
27 february 2012
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?
agenda2
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

slide32
F2(g) + 2ClO2(g) 2FClO2(g)

1

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 [F2][ClO2]

slide33
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)

slide34
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)

rate

k =

2.2 x 10-4 M/s

=

[S2O82-][I-]

(0.08 M)(0.034 M)

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)

= 0.08/M•s

practice problems1
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

slide36
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)
slide37
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
slide38
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
slide39
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
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:
28 february 2012
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?
agenda3
Agenda
  • Homework Quiz
  • 1st order reactions graphically
  • Half life calculations

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

first order overall reactions
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:

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

slide44
A product

rate

=

[A]

M/s

D[A]

-

M

= k [A]

Dt

[A] = [A]0e−kt

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

D[A]

rate = -

Dt

First-Order Reactions

rate = k [A]

= 1/s or s-1

k =

slide45
Graphical Determination of k

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

a non graphical example
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 ?
slide47
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 ?

0.88 M

ln

0.14 M

=

2.8 x 10-2 s-1

ln

ln[A]0 – ln[A]

=

k

k

[A]0

[A]

[A]0 = 0.88 M

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

[A] = 0.14 M

kt = ln[A]0 – ln[A]

= 66 s

t =

slide48
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?
29 february 2012
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?

slide50
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.

slide51
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
slide52
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:

slide53
[A]0

ln

[A]0/2

0.693

=

=

=

k

k

ln 2

k

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

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?

slide54
[A]0

ln

[A]0/2

0.693

=

=

=

=

k

k

ln 2

ln 2

0.693

=

k

k

5.7 x 10-4 s-1

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

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)

slide55
A product

# of

half-lives

[A] = [A]0/n

First-order reaction

1

2

2

4

3

8

4

16

slide56
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.

slide58
A product

rate

=

[A]2

M/s

D[A]

1

1

-

M2

= k [A]2

=

+ kt

Dt

[A]

[A]0

t½ =

D[A]

rate = -

Dt

1

k[A]0

Second-Order Reactions

rate = k [A]2

= 1/M•s

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

slide59
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.

  • 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.
slide60
The reaction 2A → B is second order with a rate constant of 51/M·min at 24oC.
  • Starting with [A]o = 0.0092 M, how long will it take for [A]t = 3.7x10-3 M?
  • Calculate the half-life of the reaction.
1 march 2012
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.

  • Starting with [A]o = 0.0092 M, how long will it take for [A]t = 3.7x10-3 M?
  • Calculate the half-life of the reaction.
agenda4
Agenda
  • Homework Quiz
  • Questions?
  • Kinetics Quiz
  • Activation Energy

Homework: p.

slide63
A product

rate

[A]0

D[A]

-

= k

Dt

[A]0

t½ =

D[A]

2k

rate = -

Dt

Zero-Order Reactions

rate = k [A]0 = k

= M/s

k =

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

[A] = [A]0 - kt

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

t½ = t when [A] = [A]0/2

slide64
Concentration-Time Equation

Order

Rate Law

Half-Life

1

1

=

+ kt

[A]

[A]0

=

[A]0

t½ =

t½ =

ln 2

2k

k

1

k[A]0

Summary of the Kinetics of Zero-Order, First-Order

and Second-Order Reactions

[A] = [A]0 - kt

rate = k

0

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

1

rate = k [A]

2

rate = k [A]2

activation energy and temperature dependence of rate constants
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
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
question
Question
  • Explain in terms of collision theory why temperature affects rate of reaction.
so when does the reaction happen
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.
slide69
+

A + B AB C + D

+

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

rate constant is temp dependent
Rate Constant is Temp. Dependent

Arrhenius equation

Eais 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
Alternate Arrhenius Equation
  • To relate k at two temperatures, T1 and T2:
slide72
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 )

determining graphically
Determining Graphically
  • slope = -2.19x104
  • slope =
determining activation energy
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.

5 march 2012
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!
agenda5
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.

with one partner
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…
7 march 2012
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.
agenda6
Agenda
  • Objective and agenda
  • Correct and explain answers to diagnostic test free response questions
with your group
With your group…
  • 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
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
agenda7
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

8 march 2012
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.)
agenda8
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.

gallery walk
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.
using the alternate arrhenius equation
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?
using the arrhenius equation
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.
frequency of collisions and orientation factor
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.
slide89
Importance of Molecular Orientation

effective collision

ineffective collision

reaction mechanisms
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.
slide91
2NO (g) + O2 (g) 2NO2 (g)

Elementary step:

NO + NO N2O2

+

Elementary step:

N2O2 + O2 2NO2

Overall reaction:

2NO + O2 2NO2

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!

13 march 2012
13 March 2012
  • Objective: You will be able to
    • identify overall reactions, intermediates and rate laws for reaction mechanisms.
agenda9
Agenda
  • Objectives and Agenda
  • Review: Reaction mechanisms
  • Elementary step examples
  • Catalysts

Homework: p. 605 #44, 45, 49, 51, 52, 54, 55, 56, 61: Tues.

slide95
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
slide96
Unimolecular reaction

Bimolecular reaction

Bimolecular reaction

A + B products

A + A products

A products

Rate Laws and Elementary Steps

rate = k [A]

rate = k [A][B]

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 sloweststep in the sequence of steps leading to product formation.

slide97
Step 1:

Step 2:

NO2 + NO2 NO + NO3

NO3 + CO NO2 + CO2

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:

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?

slide98
Step 1:

Step 2:

NO2 + NO2 NO + NO3

NO2+ CO NO + CO2

NO3 + CO NO2 + CO2

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:

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 = 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
  • rate determining step: the slowest step in the sequence of steps leading to product formation.
problem1
Problem
  • Propose a mechanism for the overall reaction:

2A + 2B → A2B2

example3
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].

  • Write the equation for the overall reaction.
  • Identify the intermediates.
  • What can you say about the relative rates of steps 1 and 2?
slide102
NO2 + F2 → NO2F + F

NO2 + F → NO2F

  • 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?
slide103
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
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-

slide105
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.
slide106
Ea

k

Ea< Ea

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

Catalyzed

Uncatalyzed

ratecatalyzed > rateuncatalyzed

catalysts
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!

week of march 12
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 = k[HBr][O2]?
13 march 20121
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

catalyst example ozone cycle
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
chlorofluorocarbons and ozone
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)
slide112
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
slide113
N2 (g) + 3H2 (g) 2NH3 (g)

Fe/Al2O3/K2O

catalyst

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.

slide114
4NH3(g) + 5O2(g) 4NO (g) + 6H2O (g)

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

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

Pt-Rh catalysts used

in Ostwald process

Ostwald Process

Pt catalyst

slide115
catalytic

CO + Unburned Hydrocarbons + O2

CO2 + H2O

converter

catalytic

2NO + 2NO2

2N2 + 3O2

converter

Catalytic Converters

slide116
Enzyme Catalysis

biological catalysts

14 march 2012
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
ap exam
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.
homework
Homework
  • Pre-lab: due tomorrow
  • Lab procedure: read by tomorrow
  • Problem set: due Monday
  • Kinetics test: Tuesday
expectations
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!
15 march 2012
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.
homework1
Homework
  • Problem Set due Monday
  • Kinetics Unit Test Tuesday
  • Gas Unit revisions due tomorrow
logistics
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!
changes to the procedure
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!
slide126
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
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.
reagents and equipment
Reagents and Equipment
  • Leave reagents at the front table. Bring your labeled pipettes to the table to fill them.
slide129
Data
  • Record your data immediately and carefully in tables in your lab notebook.
19 march 2012
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
homework2
Homework
  • Problem Set due today
  • Kinetics Test tomorrow
    • 10 MC
    • 1-2 FRQ
22 march 2012
22 March 2012
  • Objective: You will be able to:
    • determine the rate law, reaction constant and activation energy for the iodine clock reaction.
agenda10
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 8:00 am TOMORROW!!!

water baths
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.
safety
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.
cleanup
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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