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Chapter 15 - Chemical Kinetics

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Chapter 15 - Chemical Kinetics Objectives: Determine rates of reactions from graphs of concentration vs. time. Recall the conditions which affect the rates. – PowerPoint PPT presentation

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Title: Chapter 15 - Chemical Kinetics


1
Chapter 15 - Chemical Kinetics
  • Objectives
  • Determine rates of reactions from graphs of
    concentration vs. time.
  • Recall the conditions which affect the rates.
  • Recognize the order of reaction, give the rate
    equation, calculate the rate constant.
  • Use integrated rate laws and half-life equations
    in calculations.
  • Draw energy diagrams, find activation energy.
  • Identify catalysts and their properties.
  • Identify reaction intermediates.
  • Recognize the rate equation given a mechanism,
    and given the rate equation determine the
    mechanism.

2
KINETICS the study of REACTION RATES
H2O2 decomposition in an insect
H2O2 decomposition catalyzed by MnO2
Chemical engineering, enzymology, environmental
engineering, etc.
3
Kinetics and Mechanisms
  • KINETICS the study of REACTION RATES and their
    relation to the way the reaction proceeds, i.e.,
    its MECHANISM.
  • The reaction mechanism is our goal!
  • The sequence of events at the molecular level
    that control the speed and outcome of a reaction.

4
Br from biomass burning destroys stratospheric
ozone. (See R.J. Cicerone, Science, volume 263,
page 1243, 1994.)
  • Step 1 Br O3 ---gt BrO O2
  • Step 2 Cl O3 ---gt ClO O2
  • Step 3 BrO ClO light ---gt Br Cl O2
  • NET 2 O3 ---gt 3 O2
  • Identify the intermediate
  • Identify the catalyst(s)

5
Reaction Rates
  • Reaction rate change in concentration of a
    reactant or product with time.
  • Three types of rates
  • initial rate
  • average rate
  • instantaneous rate

6
Reaction Rates can be determined from a Plot
  • Blue dye is oxidized with bleach.
  • Its concentration decreases with time.
  • The rate the change in dye conc with time
    can be determined from the plot.
  • The initial rate (in this case over the first
    minute) is calculated from a tangent line
    crossing the initial concentration. Then the
    slope of the line is determined

7
Reaction Rates can be determined from a Plot
N2O5 ? NO2 O2
2
4
  • The average rate is calculated from a time
    interval.
  • The instantaneous rate is calculated at a single
    point in time (or given concentration) by drawing
    a tangent line crossing the point. Then the slope
    of the line is determined.
  • Compare average rates at the beginning and
    end of reaction.

8
Factors affecting the Rates
  • Concentrations
  • Physical State of Reactants and Products
  • Surface area
  • Temperature
  • Catalysts

Review Exp. 1 Factors affecting reactions rates
9
Concentration
Mg(s) 2 HCl(aq) ---gt
MgCl2(aq) H2(g)
0.3 M HCl
6 M HCl
  • Increasing the concentration of reactants
    _______________ the rate of the reaction.

10
Surface Area
  • Increasing the surface area of reactants
    _____________ the reaction rate.

11
Temperature
Bleach at 54 C
Bleach at 22 C
  • Increasing the temperature ____________ the rate
    of the reaction.

12
Catalysts
MnO2
2 H2O2 ---- gt 2 H2O O2
  • A _________ is present at the beginning and at
    the end of the reaction and it does not change
    but it _______________ the rate of the reaction.

13
Factors affecting Reaction Rates
  • Iodine clock reaction

1. Iodide is oxidized to iodine H2O2 2
I- 2 H -----gt 2 H2O I2 2. I2
reduced to I- with vitamin C I2 C6H8O6
----gt C6H6O6 2 H 2 I- When all vitamin
C is depleted, the I2 interacts with starch to
give a blue complex.
14
Factors affecting Reaction Rates
15
Concentration and Rate
  • To postulate a mechanism we study
  • - The reaction rate
  • and its
  • - Concentration dependence.
  • Generate a Rate Law Equation.

16
Rate Laws
  • In general for
  • a A b B --gt x X
  • Rate k AmBn
  • The exponents m, n
  • are the _______________
  • can be 0, 1, 2 or fractions
  • must be determined by ____________!

With a catalyst C
Cp
17
Interpreting Rate Laws
  • Rate k AmBnCp
  • If m 1, rxn. is 1st order in A
  • Rate k A1
  • If A doubles, then rate goes up by factor of
    __________
  • If m 2, rxn. is 2nd order in A.
  • Rate k A2
  • Doubling A increases rate by _____________
  • If m 0, rxn. is zero order.
  • Rate k A0
  • If A doubles, rate ____________________

18
Deriving Rate Laws
  • Derive rate law and k for
  • CH3CHO(g) --gt CH4(g) CO(g)
  • from experimental data for rate of disappearance
    of CH3CHO

Expt. CH3CHO Disappear of CH3CHO
(mol/L) (mol/Lsec) 1 0.10 0.020 2 0.2
0 0.081 3 0.30 0.182 4 0.40 0.318
Rate k CH3CHOn
4 Rate k 2 CH3CHOn
n 2
Look at exp 1 and 2 concentration doubles rate
cuadruples
19
Deriving Rate Laws
  • Rate of rxn
  • Here the rate goes up by ________ when initial
    concentration doubles. Therefore, we say this
    reaction is __________ order.
  • Now determine the value of k. Use any exp.
    Data
  • Using k you can calculate rate at other values
    of CH3CHO at same T.

20
Concentration and Time
  • What is the concentration of reactant as a
    function of time?
  • Consider First Order reactions

Integrating we get
21
Integrated First Order Law
A / A0 fraction remaining after time t has
elapsed.
22
The decomposition of a certain insecticide in
water follows first-order kinetics with a rate
constant of 1.45 year-1 at 12oC. A quantity of
this insecticide is washed into a lake on June 1,
leading to a concentration of 5.0 x 10-7 g/cm3.
Assume that the average temperature of the lake
is 12oC.
  • a) What is the concentration of the insecticide
    on June 1 of the following year?b) How long will
    it take for the concentration of the insecticide
    to drop to 3.0 x 10-7 g/cm3?

23
Using Integrated Rate Laws
  • All 1st order reactions have straight line plot
    for ln A vs. time.
  • And 2nd order gives straight line for plot of
    1/A vs. time.

24
Using Integrated Rate Laws
  • In an experiment for
  • 2 N2O5(g) ---gt 4 NO2(g) O2(g)
  • Time (min) N2O50 (M)
  • 0 1.00
  • 1.0 0.705
  • 2.0 0.497
  • 5.0 0.173

If it were zero order
Data of conc. vs. time plot do not fit straight
line.
25
Using Integrated Rate Laws
  • In an experiment for
  • 2 N2O5(g) ---gt 4 NO2(g) O2(g)
  • Time (min) N2O50 (M)
  • 0 1.00
  • 1.0 0.705
  • 2.0 0.497
  • 5.0 0.173

If it were first order
ln N2O50 0 -0.35 -0.70 -1.75
Plot of ln N2O5 vs. time is a straight line!
Calculate the ln
26
Using Integrated Rate Laws
Plot of ln N2O5 vs. time is a straight line!
Eqn. for straight line y mx b
27
The gas phase decomposition of hydrogen peroxide
at 400 oC is second order in H2O2. In one
experiment, when the initial concentration of
H2O2 was 0.246 M, the concentration of H2O2
dropped to 3.39 x 10-2 M after 25.9 seconds had
passed. What is the rate constant for the
reaction?
  • 2 H2O2 ? 2 H2O O2

28
Half-Life
  • HALF-LIFE is the time it takes for 1/2 a sample
    is disappear. For 1st order reactions, the
    concept of HALF-LIFE is especially useful.

29
Half-Life
  • Reaction is 1st order decomposition of H2O2.
  • Reaction after 1 half-life.
  • 1/2 of the reactant has been consumed and 1/2
    remains.

30
Half-Life
  • After 2 half-lives _____ of the reactant remains.
  • After 3 half-lives ____ of the reactant remains.

31
Half-Life
ln R / R0 k t
  • A / A0 fraction remaining
  • when t t1/2 then fraction remaining _________
  • ln (____) - k t1/2

32
In an experiment, it is determined that 75 of a
sample of HCO2H (formic acid) has decomposed in
72 seconds following first-order kinetics.
Determine t1/2 for this reaction. HCO2H ? CO2
H2
ln R / R0 k t
t1/2 0.693 / k
33
Mechanisms
  • Mechanism how reactants are converted to
    products at the molecular level.

RATE LAW ----gt MECHANISM experiment
----gt theory
34
Activation Energy
  • Molecules need a minimum amount of energy to
    react. Visualized as an energy barrier -
    activation energy, Ea.

Reaction coordinate diagram
35
Activation Energy
  • Conversion of cis to trans-2-butene requires
    twisting around the CC bond.
  • Rate k trans-2-butene

36
Transition State
Cis
Trans
Transition state
37
Mechanisms
  • Reaction passes thru a TRANSITION STATE where
    there is an activated complex that has sufficient
    energy to become a product.

ACTIVATION ENERGY, Ea energy reqd to form
activated complex. Here Ea ___________
38
Mechanisms
  • Also note that trans-butene is MORE STABLE than
    cis-butene by about 4 kJ/mol.
  • Therefore, cis ---gt trans is _________________
  • This is the connection between thermo-dynamics
    and kinetics.

39
Effect of Temperature
In ice at 0 oC
  • Reactions generally occur slower at lower T.

Room temperature
Iodine clock reaction, book page 705. H2O2 2
I- 2 H --gt 2 H2O I2
40
Activation Energy and Temperature
In general, differences in activation energy
cause reactions to vary from fast to slow.
  • Reactions are __________ at a higher T because
    a larger fraction of reactant molecules have
    enough energy to convert to product molecules.

41
Collision Theory
  • Molecules must collide with one another
  • Molecules must collide with sufficient ________
    to break bonds
  • Molecules must collide in an orientation that can
    lead to rearrangement of the atoms.

42
Arrhenius Equation
Frequency factor related to frequency of
collisions with correct geometry.
Plot ln k vs. 1/T ---gt straight line. slope
-Ea / R
43
Collision Theory
  • Reactions require
  • (a) activation energy and
  • (b) correct geometry.
  • O3(g) NO(g) ---gt O2(g) NO2(g)

2. Activation energy and geometry
1. Activation energy
44
Mechanisms
  • Most reactions involve a sequence of elementary
    steps.
  • 2 I- H2O2 2 H ---gt I2 2 H2O
  • Rate k I- H2O2
  • NOTE
  • 1. Rate law comes from experiment.
  • 2. Order and stoichiometric coefficients not
    necessarily the same!
  • 3. Rate law reflects all chemistry down to and
    including the slowest step in multistep reaction.

45
Mechanisms
Most reactions involve a sequence of elementary
steps. 2 I- H2O2 2 H ---gt I2 2
H2O Rate k I- H2O2
Proposed Mechanism Step 1 slow HOOH I- --gt
HOI OH- Step 2 fast HOI I- --gt I2
OH- Step 3 fast 2 OH- 2 H --gt 2
H2O Rate of the reaction controlled by slow step
RATE DETERMINING STEP, Rate can be no faster
than RDS!
46
Mechanisms
Most reactions involve a sequence of elementary
steps. 2 I- H2O2 2 H ---gt I2 2
H2O Rate k I- H2O2
Proposed Mechanism Step 1 slow HOOH I- --gt
HOI OH- Step 2 fast HOI I- --gt I2
OH- Step 3 fast 2 OH- 2 H --gt 2 H2O
  • Elementary Step 1 is bimolecular and involves I-
    and HOOH. Therefore, this predicts the rate law
    should be
  • Rate a I- H2O2 as observed!!
  • The species HOI and OH- are ________________.

47
Catalysts and Activation Energy
MnO2 catalyzes decomposition of H2O2 2 H2O2 ---gt
2 H2O O2
48
Catalysts and Activation Energy
  • Iodine-Catalyzed Isomerization of Cis-2-Butene

49
Remember
  • Go over all the contents of your textbook.
  • Practice with examples and with problems at the
    end of the chapter.
  • Practice with OWL tutor.
  • Practice with the quiz on CD of Chemistry Now.
  • Work on your OWL assignment for Chapter 15.
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