Chapter 16 Kinetics

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Chapter 16Kinetics
Dinosaurs generated enough heat to sustain its
biochemical reactions at high rates.
Reaction rate f(temperature)
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What is chemical kinetics?
Deals with the speed (rate) of a chemical
reaction and its reaction mechanism.
Describes the change in concentration as a
function of time.
Quantitatively
Qualitatively
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Section 16.1 Factors that influence reaction
rate
Each specific reaction has its own characteristic
reaction rate.
We can control 4 factors that affect the rate of
a given reaction Concentration of
reactants Physical state of reactants
Temperature of reaction Presence of a catalyst
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Section 16.1 Factors that influence reaction
rate
(1) Concentration Molecules must collide in
order to react.
The reaction rate changes throughout the course
of a reaction.
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Section 16.1 Factors that influence reaction
rate
(2) Physical state of reactants Molecules must
collide in order to react.
Reactants in same physical state ? random thermal
motion brings them into contact Reactants in
different physical states ? contact between
reactants occurs only at the

interface between the phases
Example reactants orange blue
The more finely divided a solid or
liquid reactant, the greater its surface area
per unit volume ? More contact with
other reactants ? Faster reaction.
solid aqueous (interface only)
aqueous aqueous
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Section 16.1 Factors that influence reaction
rate
(3) Temperature of reaction Molecules must
collide with enough energy to react.
Two aspects to this (1) At higher temperatures,
more collisions occur at a given time. (2) At
higher temperatures, the energy of collisions is
higher (K.E. of molecules is higher).
Example refrigeration versus cooking
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Section 16.2 Expressing reaction rate
quantitatively
Rate a change in some variable per unit of time
Analogy with speed
Reaction A ? B
reaction rate the changes in concentrations of
reactants or products per unit time
Reactant concentrations decrease, while product
concentrations increase.
Reaction A ? B
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Section 16.2 Expressing reaction rate
quantitatively
In most reactions, not only the concentration
changes, but the reaction rate also
changes. Therefore, we can define three reaction
rates Average reaction rate how fast the
concentration changes of the entire time
period Instantaneous reaction rate the
reaction rate at an instant in time Initial
reaction rate the instantaneous rate at the
moment the reactants are mixed
Example Reaction involved in the formation of
photochemical smog (ethylene ozone)
Reaction rate (rate of decrease of reactants)
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Section 16.2 Expressing reaction rate
quantitatively
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Section 16.2 Expressing reaction rate
quantitatively
Rates for reactants and products
General formula (a, b, c, d ?
coefficients) General equation
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Section 16.2 Expressing reaction rate
quantitatively
General formula (a, b, c, d ?
coefficients) General equation

• Express the rate in terms of changes in H2,
  O2, and H2O with time.
• When O2 is decreasing at 0.23 mol/L sec, at
  what rate is H2O increasing?

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Section 16.3 Rate laws
Experimentally determined ? not determined from
reaction stoichiometry
General reaction aA bB ? cC dD
Rate law rate kAmBn
where A and B are concentrations of
reactants A and B k is the rate law
constant specific for a given reaction at a
given temperature m and n are the
reaction orders defines how the rate is
affected by reactant

concentration
More on reaction orders m and n rate change /
concentration change Examples If the rate
doubles when A doubles ? A1 and m 1 If the
rate quadruples when B doubles ? B2 and n
2 If the rate does not change when A doubles ?
A0 and m 0
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Section 16.3 Rate laws
All terms in the rate law (rate kAmBn.)
must be determined experimentally. A and B
measured and rate, rate constant (k), reaction
orders (m, n) are deduced from these measurements.

• Measuring rates many methods
• Conductometric methods used when a nonionic
  reactant forms ionic products
• Example (CH3)3CBr (l) H2O (l) ?
  (CH3)3COH (l) H (aq) Br- (aq)

(2) Manometric methods used when a reaction
involves a change in the number of moles of
a gaseous reactant or product ? reaction rate
determined by the change in pressure over
time Example Zn (s) 2 CH3COOH (aq) ?
Zn2 (aq) 2 CH3COO- (aq) H2 (g)
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Section 16.3 Rate laws
(3) Spectrometric methods used when one of the
reactants of products absorbs (or emits)
certain wavelengths of light
Example NO (g, colorless) 2 O3 (g, colorless)
? O2 (g, colorless) NO2 (g, brown)
Lightin
Lightout
NO2
Transparent Reaction Cell
Light Source
Light Detector
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Section 16.3 Rate laws
(4) Direct chemical methods used for reactions
that can be easily slowed or stopped
Example Measure respiration rate by killing
bacteria with HgCl2 to stop respiration.
Respiration O2 CH2O ? CO2 energy
BOD bottle (Biological Oxygen Demand)
Measure O2 concentration at t0 and t24
hrs. Change in O2 concentration
O2 (t0) O2 (t24 hrs)
Respiration rate change in O2 concentration
time (24
hours)
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Section 16.3 Rate laws
Determining Reaction Order
First, some terminologyindividual reaction
order vs. overall reaction order
Example 2 NO (g) 2 H2 (g) ? N2 (g) 2 H2O (g)
Rate kNO2H2 Individual Reaction
is second order with respect to NO and first
order with respect to H2. Overall Reaction is
third order overall (sum of individual reaction
orders).
Rate kNOO3 Rate k(CH3)3CBrH2O0 Rat
e kCHCl3Cl21/2
A zero order reaction order means that the
reaction does not depend on the concentration of
that reactant.
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Section 16.3 Rate laws
Determining Reaction Order
We have been determining reaction orders from a
known rate law (i.e. Rate kNOO3) When the
rate law is not known, use data from a series of
experiments with different reactant
concentrations to determine initial reaction
rates.
Change one reactant concentration, while keeping
the other constant.
Reaction is what order with respect to O2? With
respect to NO? Overall?
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Section 16.3 Rate laws
Determining Reaction Order
We have been determining reaction orders from a
known rate law (i.e. Rate kNOO3) When the
rate law is not known, use data from a series of
experiments with different reactant
concentrations to determine initial reaction
rates.
Change one reactant concentration, while keeping
the other constant.
Reaction is what order with respect to O2?
Double O2, double reaction rate 1st order w.r.t
O2 reaction order rate change / concentration
change
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Section 16.3 Rate laws
Determining Reaction Order
We have been determining reaction orders from a
known rate law (i.e. Rate kNOO3) When the
rate law is not known, use data from a series of
experiments with different reactant
concentrations to determine initial reaction
rates.
Change one reactant concentration, while keeping
the other constant.
Reaction is what order with respect to NO?
Double NO, quadruple reaction rate 2nd order
w.r.t NO
Rate law rate kO2NO2 Overall reaction
order 3rd order
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Section 16.3 Rate laws
Determining the Rate Constant
Simply, solve for k.
Rate law rate kO2NO2
What is k for this reaction?
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Section 16.4 Integrated rate laws
So far, we have not considered the time factor in
the rate law equations.
Rate kO2NO2 This equation says what the
rate will be when O2 is X and NO is Y, but
does not tell use how long it will take for X
moles of NO to be used up (for example).
Integrated rate laws consider the time factor
and are derived from equations
we have already seen using
calculus
For reaction A ? B
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Section 16.4 Integrated rate laws
Example At 1000 ºC, cyclobutane (C4H8)
decomposes in a first-order reaction, with The
very high rate constant of 87 s-1, to two
molecules of ethylene (C2H4). If the initial
cyclobutane concentration is 2.00 M, what is the
concentration after 0.010 s? What fraction of
cyclobutane has decomposed in this time?
At 25 ºC, hydrogen iodide breaks down very slowly
to hydrogen and iodine rate kHI2 The rate
constant at 25 ºC is 2.4 x 10-21 L/mol sec. If
0.0100 mol of HI(g) is placed in a 1.0 L
container, how long will it take for the
concentration of HI to reach 0.00900 mol/L?
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Suggested Problems
16.2, 16.3, 16.12, 16.14, 16.16, 16.20, 16.26,
16.38, 16.34