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CHEMICAL KINETICS CHAPTER 13

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CHEMICAL KINETICS CHAPTER 13 VI. Temperature and Rxn Rate A. Nanoscale Explanation as to why increasing temperature increase reaction rate. B. – PowerPoint PPT presentation

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Title: CHEMICAL KINETICS CHAPTER 13


1
CHEMICAL KINETICSCHAPTER 13
2
I. Introduction
  • A. Definition of Chemical Kinetics
  • The study of the speed or rate of reactions and
    the nanoscale pathways or processes by which
    reactants are transformed into products.
  • B. Examples of Reactions and Rates
  • Rusting of Iron
  • Combustion Reaction
  • C. Significance of Studying Kinetics

3
  • D. Factors Affecting Reaction Rate
  • 1. Concentration of Reactants
  • 2. Temperature
  • 3. Presence of a Catalyst
  • 4. Surface Area of a Solid Reactant or Catalyst
  • 5. Properties of Reactants and Products

4
II. Understanding Reaction Rates
  • A. Kinetic Molecular Theory
  • Matter composed of particles in constant motion.
  • Increase in temperature increases particles
    kinetic energy.
  • B. Collision Theory
  • For a reaction to occur , reactant molecules
    must collide with the proper orientation and with
    an energy greater than some minimum value.
  • Activation Energy (Ea) minimum energy
    required for reaction to occur.

5
Importance of Orientation
One hydrogen atom can approach another from any
direction
Effective collision the I atom can bond to the C
atom to form CH3I
and reaction will still occur the spherical
symmetry of the atoms means that orientation does
not matter.
Ineffective collision orientation is important
in this reaction.
6
Distribution of Kinetic Energies
At higher temperature (red), more molecules have
the necessary activation energy.
7
  • C. Transition State Theory
  • 1. Note Reaction Profile

CO(g) NO2(g) ?
CO2(g) NO(g)
8
  • 2. Note Ea for either forward or
    reverse reaction.
  • At a given temperature, the
    higher the energy barrier, the slower the
    reaction.
  • 3. Transition State or Activated Complex
  • Transition structure (between reactants
    and
  • products) which is always found
    on top of the
  • energy hill (energy of
    activation).
  • 4. Is reaction, exothermic or endothermic
    as written from left to right?

9
  • Reconsideration of Factors Affecting Reaction
    Rate!!
  • 1. Concentration
  • 2. Temperature
  • 3. Catalyst

10
III. Rates of Reactions
  • A. Definition
  • 1. Reaction rate expresses how much
    product is appearing or how much reactant in
    disappearing per unit time.
  • 2. Units for reaction rates
  • (Examples)
  • Ms-1 or M/s or
    mol L-1s-1

11
For Reaction A ? P

  • Rate of Disappearance of A - ?A / ?t
  • Rate of Formation of P ?P / ?t

12
  • B. Example Average Rate Determination
  • For Rxn of Cisplatin and Water
  • H2O Pt(NH3)2Cl2 ? Pt(NH3)2Cl(H2O)
    Cl-
  • Cisplatin
  • Time (min) Cisplatin
    (mol/L)
  • 0 0.01000
  • 200 0.00747
  • 400
    0.00558
  • What is average rate of
    disappearance of cisplatin
  • (in mol L-1 min-1) for the first
    200 min?
  • What is average rate of
    disappearance of cisplatin
  • (in mol L-1 min-1) for the next 200
    min?

13
  • C. Instantaneous Rate Determination
  • Significance of Measuring Instantaneous
    Rates!!

14
  • D. Reaction Rates and Stoichiometry
  • Given the reaction
  • 2 N2O5(g) ? 4 NO2(g)
    O2 (g)
  • If the rate of NO2 formation is
    0.060 mol L-1 s-1
  • 1. What is the rate of disappearance of
    N2O5?
  • 2. What is the rate of formation of O2?

15
IV. Concentration and Rxn Rate
  • A. Rate Law Equation
  • 1. An equation that relates the rate of a
  • reaction to the concentrations of
    reactants
  • (and catalyst) raised to various
    powers.
  • 2. Must be experimentally determined!!

16
  • 3. For reaction
  • A B ? C
  • Rate is proportional to reactant
    concentrations
  • Rate kAm Bn k
    rate constant
  • (exponents m and n must be
    experimentally determined).
  • 4. For reaction
  • 2 NO2(g) F2(g) ? 2 NO2F(g)
  • (experimentally determined Rate Law
    is)
  • Rate k NO21F21 k NO2F2
  • Exponents not necessarily same as rxn
    coefficients!!

17
  • 4. For hypothetical reaction
  • 2 A(g) B2(g) ? 2 AB(g)
  • (experimentally determined Rate Law
    is)
  • Rate k A2
  • Not all reactants necessarily show up
    in the rate law equation!!

18
  • B. Reaction Order
  • For the general equation aA bB
    ? pP
  • The rate equation is
  • Rate k Am Bn
  • m and n are experimentally determined and are
    usually integers (0, 1, 2, 3, ). They may be
    fractions.

19
  • This reaction is said to m th order with respect
    to A and n th order with respect
  • to B.
  • The overall reaction order is the sum of the
    individual orders, or
  • Overall Reaction Order m n

20
  • Example
  • For the following reaction
  • 2 NO(g) Cl2(g) ? 2 NOCl(g)
  • The observed rate law is
  • Rate k NO2 Cl2
  • What is the reaction order with respect to NO?
    What is the reaction order with respect to Cl2?
    What is the overall reaction order?
  • How would the rate of the reaction be affected
    by doubling the concentration of both NO and
    Cl2?

21
  • C. Determination of Rate Law Exponents
  • Done experimentally by measuring initial
    rates for several different known concentrations
    of reactants.
  • Consider the reaction
  • 2 NO(g) 2 H2(g) ? N2(g) 2 H2O(g)
  • Given the information on the next slide
  • 1. Determine the rate law.
  • 2. What is the order of the reaction?
  • 3. What is the value of the rate constant?
    units

22
  • Initial Concentration
    (M) Rate
  • Experiment NO H2
    mol / L.s

1 0.100 0.100 1.23 x 10-3
2 0.100 0.200 2.45 x 10-3
3 0.200 0.100 4.93 x 10-3
4 0.300 0.100 1.11 x 10-2
23
  • 1. Determine the rate law.
  • a. Determine general form of rate law.
  • rate k NOmH2n
  • b. Determine exponents.
  • For each reactant, compare two
    experiments or trials where its
    concentration is changing and all other
    reactant concentrations are held
    constant.

24
  • Methods For Determining Rate Law Exponents
  • Method 1 solve analytically
  • Substitute data into rate law and
    compare
  • Divide equation with larger rate
    by eq. with
  • smaller rate. Cancel terms and
    solve.
  • or simplifying
  • Method 2 - solve by inspection
  • How does changing conc. affect
    rate?

25
  • 2. Determine order of the reaction.
  • 3. Determine rate constant (include units).
  • Use rate law and either set of data.
  • 4. What is the rate of the reaction when
  • NO H2 0.200 M ?

26
V. Integrated Rate Law
  • Equations derived from rate law (by using
    calculus) which are convenient for solving
    concentration versus time problems.
  • For
  • 1. First Order Rxns only one covered
  • 2. Second Order Rxns
  • 3. Zero Order Rxns

27
  • A. Integrated First Order Rate Law
  • For reaction
  • aA ? Product
  • rate - ?A / ?t kA
  • and using calculus
  • where Ao is the concentration of A at
    time zero (t 0) and At is the
    concentration at time t.
  • A and A0 may be replaced by quantities
    that are proportional to concentration !!!!

28
  • B. Problem
  • The sugar, sucrose, will undergo the
    following
  • (first order) hydrolysis reaction
  • C12H22O11 H2O ? C6H12O6 C6H12O6
  • sucrose
    glucose fructose
  • with a rate constant of 6.2 x 10-5 s-1
    at 35oC.
  • A sample of 0.20 mol of sucrose was
    initially
  • dissolved in a total volume of 500 mL.
  • 1. What is the sucrose conc. after 2
    hours?

29
  • 2. What will be the glucose concentration
    after the 2 hours have elapsed?
  • 3. How many minutes will it take for the
  • sucrose concentration to drop to
    0.30M?

30
  • C. Half-Life and First Order Reactions
  • 1. Definition (t1/2) the time required
    for the concentration of a reactant to fall to
    one half its initial value.
  • 2. Significance
  • Useful in describing radioactive decay
    rates
  • Useful in describing rates of 1st order
    reactions
  • 3. Equation derived from Integrated Rate
    Law

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32
  • 4. Problem - Radioactive Iodine-131 has t1/2
    of 8
  • days. If you had a sample of 10,000
    I-131
  • atoms initially, how many I-131 atoms would
  • remain after 32 days?

33
  • 5. Problem The first order hydrolysis
    reaction of
  • sucrose
  • C12H22O11 H2O ? C6H12O6 C6H12O6
  • sucrose
    glucose fructose
  • has a rate constant of 6.17 x 10-4 s-1
    under
  • experimental conditions.
  • a. What is the half-life for the hydrolysis of
  • sucrose?
  • b. How many minutes are required for
    75 of the initial sucrose to react?

34
VI. Temperature and Rxn Rate
  • A. Nanoscale Explanation as to why increasing
    temperature increase reaction rate.
  • B. Mathematical Relationship
  • Arrhenius Equation
  • (not responsible for problem solving)

35
VII. Reaction Mechanisms
  • A. Introduction
  • Reaction mechanism is a series of
    elementary reactions or simple steps whose
    overall effect is given by the net chemical
    reaction (equation).
  • 1. It is a theory of how the reaction
    occurs which is based on experimental data.
  • 2. Cannot be absolutely proven.
  • 3. Steps must be elementary reactions.

36
  • B. Elementary Reaction
  • 1. Definition- the simplest step in what
    is
  • often a multi-step mechanism for an
  • observed chemical reaction.
  • a. The equation for an elementary
    reaction shows exactly which molecules,
    atoms, or ions take part in the
    elementary reaction.
  • b. For an elementary reaction, the
    rate law is directly determined from the
    elementary reaction.

37
  • 2. Elementary Reactions in Mechanisms
  • Types
  • a. Unimolecular Reaction structure of a
    single particle (atom, molecule, or ion)
    rearranges to produce a different particle
    or particles.
  • b. Bimolecular Reaction - two particles
    (atoms, ions, or molecules) collide and
    rearrange into products.
  • c. Termolecular Reaction (less likely)

38
  • 3. Problems
  • Identify the type of elementary reaction and
    give the rate law for the following elementary
  • reactions
  • a. Cl Cl ? Cl2
  • b. N2O5 ? NO2 NO3

39
  • C. Properties of Valid Mechanisms
  • 1. Must consist of only unimolecular,
    bimolecular, or termolecular elementary
    reactions. (True for any mechanisms given to
  • you.)
  • 2. Sum of the elementary reactions should be
    equal to the overall reaction equation.
  • 3. Should predict the experimentally observed
    rate law.
  • The overall rate of the reaction is
    dependent on the slowest step in the mechanism
    - the rate-limiting step.

40
  • D. Example Mechanism Problems
  • 1. For the overall reaction
  • 2 NO2Cl ? 2 NO2 Cl2
  • the following mechanism is proposed.
  • NO2Cl ? NO2 Cl (slow)
  • NO2Cl Cl ? NO2 Cl2 (fast)
  • a. Does the sum of the elementary
    processes equal the overall reaction?
  • b. What is the rate law for the overall rxn?
  • c. Identify any reaction intermediates.

41
  • 2. For the overall reaction
  • (CH3)3CCl OH- ? (CH3)3COH Cl-
  • there are two proposed mechanisms.
  • 1) Concerted mechanism
  • (CH3)3CCl OH- ? (CH3)3COH Cl-
  • 2) Two Step Mechanism
  • (CH3)3CCl ? (CH3)3C Cl -
    (slow)
  • (CH3)3C OH - ? (CH3)3COH (fast)

42
  • From kinetic data, the correct rate law for
    the overall reaction is
  • rate k(CH3)3CCl
  • Questions
  • 1. Which is the correct mechanism? Why?
  • 2. What is the order of the overall
    reaction?
  • 3. Identify reaction intermediates in each
  • proposed mechanism.

43
VIII. Catalysts
  • A. Definition
  • Catalyst a species that increases the rate of
    an overall reaction but is not consumed in the
    reaction.
  • Not shown in overall reaction.
  • Will show up in rate law for catalyzed
    reaction.

44
  • B. How Do They Increase Reaction Rate?
  • 1. Catalysts alter / participate in rxn
    mechanism.
  • 2. Lower activation energy. Speeds reaction.
  • 3. See Fig 13.17 (next slide)
  • 4. Catalyst changes kinetics, but not
  • thermodynamics of reaction.
  • Increases speed of reaction.
  • Does not change net energy of reaction, type
    of product produced, or direction of reaction.

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46
  • C. Homogeneous vs. Heterogeneous Catalysts
  • Heterogeneous Catalyst- catalyst in
    different phase from reacting substance.
  • (hydrogenation of vegetable oils next slide)
  • (catalytic converters in autos)
  • Homogeneous Catalyst catalyst in same phase
    as reacting substance.
  • (enzymes)

47
Heterogeneous Catalysis
Hydrogen is adsorbed onto the surface of a nickel
catalyst. A CC approaches
and is adsorbed.
Hydrogen atoms attach to the carbon atoms, and
the molecule is desorbed.
48
  • D. Enzymes (Homogeneous Catalysts)
  • 1. Protein that catalyzes reaction.
  • 2. Most efficient catalysts known to man.
  • 3. Specifically binds to reactants
    (substrates),
  • holding them in correct position for
    reaction to occur.
  • 4. Lower activation energy by stabilizing
  • transition state or altering
    mechanism.
  • 5. Examples Lysozyme Next slide

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