15 february 2012
This presentation is the property of its rightful owner.
Sponsored Links
1 / 137

15 February 2012 PowerPoint PPT Presentation


  • 112 Views
  • Uploaded on
  • Presentation posted in: General

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.

Download Presentation

15 February 2012

An Image/Link below is provided (as is) to download presentation

Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author.While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server.


- - - - - - - - - - - - - - - - - - - - - - - - - - E N D - - - - - - - - - - - - - - - - - - - - - - - - - -

Presentation Transcript


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.


Chemical kinetics

Chemical Kinetics


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.


15 february 2012

A B

rate =

D[A]

D[B]

rate = -

Dt

Dt


15 february 2012

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.


15 february 2012

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

time

393 nm

Detector

light

red-brown

t1< t2 < t3

D[Br2] aD Absorption


15 february 2012

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.


15 february 2012

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)


15 february 2012

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)


15 february 2012

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)?


    15 february 2012

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


    15 february 2012

    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!


    15 february 2012

    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.


    15 february 2012

    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]


    15 february 2012

    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)


    15 february 2012

    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


    15 february 2012

    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]


    15 february 2012

    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)


    15 february 2012

    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


    15 february 2012

    • 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)


    15 february 2012

    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


    15 february 2012

    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


    15 february 2012

    • 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


    15 february 2012

    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 =


    15 february 2012

    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 ?


    15 february 2012

    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 =


    15 february 2012

    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?


    15 february 2012

    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.


    15 february 2012

    • 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


    15 february 2012

    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:


    15 february 2012

    [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?


    15 february 2012

    [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)


    15 february 2012

    A product

    # of

    half-lives

    [A] = [A]0/n

    First-order reaction

    1

    2

    2

    4

    3

    8

    4

    16


    15 february 2012

    • 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.


    15 february 2012

    • Calculate the half-life of the decomposition of N2O5:

      2N2O5 4NO2(g) + O2(g)


    15 february 2012

    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


    15 february 2012

    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.


    15 february 2012

    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.


    15 february 2012

    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


    15 february 2012

    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.


    15 february 2012

    +

    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:


    15 february 2012

    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.


    15 february 2012

    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.


    15 february 2012

    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!


    15 february 2012

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

    Mechanism:


    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.


    15 february 2012

    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


    15 february 2012

    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.


    15 february 2012

    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?


    15 february 2012

    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?


    15 february 2012

    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?


    15 february 2012

    • 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-


    15 february 2012

    • 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.


    15 february 2012

    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)


    15 february 2012

    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


    15 february 2012

    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.


    15 february 2012

    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


    15 february 2012

    catalytic

    CO + Unburned Hydrocarbons + O2

    CO2 + H2O

    converter

    catalytic

    2NO + 2NO2

    2N2 + 3O2

    converter

    Catalytic Converters


    15 february 2012

    Enzyme Catalysis

    biological catalysts


    15 february 2012

    Binding of Glucose to Hexokinase


    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!


    15 february 2012

    • 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.


    15 february 2012

    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


    What s the purpose

    What’s the purpose?


    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.


  • Login