Contributed by:
The highlights are:
1. Reaction rates
2. Rate laws
3. Integrated Rate laws
4. Half-life
5. Arrhenius Equation
6. Mechanisms
1.
Chemistry, The Central Science, Chapter 14,
10th edition
Theodore L. Brown; H. Eugene LeMay, Jr.;
and Bruce E. Bursten
Chemical Kinetics
John D. Bookstaver
St. Charles Community College
St. Peters, MO
2006, Prentice Hall, Inc. Chemical
Modified by S.A. Green, 2006 Kinetics
2.
Kinetics
• Studies the rate at which a chemical
process occurs.
• Besides information about the speed at
which reactions occur, kinetics also
sheds light on the reaction mechanism
(exactly how the reaction occurs).
Chemical
Kinetics
3.
Outline: Kinetics
Reaction Rates How we measure rates.
How the rate depends on amounts
Rate Laws of reactants.
How to calc amount left or time to
Integrated Rate Laws reach a given amount.
How long it takes to react 50% of
Half-life reactants.
Arrhenius Equation How rate constant changes with T.
Link between rate and molecular
Mechanisms scale processes.
Chemical
Kinetics
4.
Factors That Affect Reaction Rates
• Concentration of Reactants
As the concentration of reactants increases, so does the likelihood that
reactant molecules will collide.
• Temperature
At higher temperatures, reactant molecules have more kinetic energy,
move faster, and collide more often and with greater energy.
• Catalysts
Speed rxn by changing
mechanism.
Chemical
Kinetics
5.
Reaction Rates
Rxn Movie
Rates of reactions can be determined by
monitoring the change in concentration of
either reactants or products as a function of
time. [A] vs t
Chemical
Kinetics
6.
Reaction Rates
C4H9Cl(aq) + H2O(l) C4H9OH(aq) + HCl(aq)
[C4H9Cl] M
In this reaction, the
concentration of
butyl chloride,
C4H9Cl, was
measured at various
times, t.
Chemical
Kinetics
7.
Reaction Rates
C4H9Cl(aq) + H2O(l) C4H9OH(aq) + HCl(aq)
Average Rate, M/s
The average rate of
the reaction over
each interval is the
change in
concentration divided
by the change in time:
Chemical
Kinetics
8.
Reaction Rates
C4H9Cl(aq) + H2O(l) C4H9OH(aq) + HCl(aq)
• Note that the average
rate decreases as the
reaction proceeds.
• This is because as the
reaction goes forward,
there are fewer
collisions between
reactant molecules.
Chemical
Kinetics
9.
Reaction Rates
C4H9Cl(aq) + H2O(l) C4H9OH(aq) + HCl(aq)
• A plot of concentration
vs. time for this reaction
yields a curve like this.
• The slope of a line
tangent to the curve at
any point is the
instantaneous rate at
that time.
Chemical
Kinetics
10.
Reaction Rates
C4H9Cl(aq) + H2O(l) C4H9OH(aq) + HCl(aq)
• The reaction slows
down with time because
the concentration of the
reactants decreases.
Chemical
Kinetics
11.
Reaction Rates and Stoichiometry
C4H9Cl(aq) + H2O(l) C4H9OH(aq) + HCl(aq)
• In this reaction, the ratio
of C4H9Cl to C4H9OH is
1:1.
• Thus, the rate of
disappearance of
C4H9Cl is the same as
the rate of appearance
of C4H9OH.
-[C4H9Cl] [C4H9OH]
Rate = = Chemical
t t Kinetics
12.
Reaction Rates and Stoichiometry
• What if the ratio is not 1:1?
H2(g) + I2(g) 2 HI(g)
• Only 1/2 HI is made for each H2 used.
Chemical
Kinetics
13.
Reaction Rates and Stoichiometry
• To generalize, for the reaction
aA + bB cC + dD
Reactants (decrease) Products (increase)
Chemical
Kinetics
14.
Concentration and Rate
Each reaction has its own equation that
gives its rate as a function of reactant
concentrations.
this is called its Rate Law
To determine the rate law we measure the rate
at different starting concentrations.
Chemical
Kinetics
15.
Concentration and Rate
Compare Experiments 1 and 2:
when [NH4+] doubles, the initial rate doubles.
Chemical
Kinetics
16.
Concentration and Rate
Likewise, compare Experiments 5 and 6:
when [NO2-] doubles, the initial rate doubles.
Chemical
Kinetics
17.
Concentration and Rate
This equation is called the rate law,
and k is the rate constant.
Chemical
Kinetics
18.
Rate Laws
• A rate law shows the relationship between the reaction
rate and the concentrations of reactants.
For gas-phase reactants use PA instead of [A].
• k is a constant that has a specific value for each reaction.
• The value of k is determined experimentally.
“Constant” is relative here-
k is unique for each rxn
k changes with T (section 14.5)
Chemical
Kinetics
19.
Rate Laws
• Exponents tell the order of the reaction with
respect to each reactant.
• This reaction is
First-order in [NH4+]
First-order in [NO2−]
• The overall reaction order can be found by
adding the exponents on the reactants in the
rate law.
• This reaction is second-order overall.
Chemical
Kinetics
20.
Integrated Rate Laws
Consider a simple 1st order rxn: A B
Differential form:
How much A is left after time t? Integrate:
Chemical
Kinetics
21.
Integrated Rate Laws
The integrated form of first order rate law:
Can be rearranged to give:
[A]0 is the initial concentration of A (t=0).
[A]t is the concentration of A at some time, t,
Chemical
during the course of the reaction. Kinetics
22.
Integrated Rate Laws
Manipulating this equation produces…
…which is in the form y = mx + b
Chemical
Kinetics
23.
First-Order Processes
If a reaction is first-order, a plot of ln [A]t
vs. t will yield a straight line with a slope
of -k.
So, use graphs to determine rxn order.
Chemical
Kinetics
24.
First-Order Processes
Consider the process in
which methyl isonitrile is
converted to acetonitrile.
CH3NC CH3CN
How do we know this
is a first order rxn?
Chemical
Kinetics
25.
First-Order Processes
CH3NC CH3CN
This data was
collected for this
reaction at 198.9°C.
for all time intervals?
Chemical
Kinetics
26.
First-Order Processes
• When ln P is plotted as a function of time, a
straight line results.
The process is first-order. Chemical
Kinetics
k is the negative slope: 5.1 10-5 s-1.
27.
Second-Order Processes
Similarly, integrating the rate law for a
process that is second-order in reactant A:
Rearrange, integrate:
also in the form y = mx + b
Chemical
Kinetics
28.
Second-Order Processes
So if a process is second-order in A, a
plot of 1/[A] vs. t will yield a straight line
with a slope of k.
First order:
If a reaction is first-order, a plot of ln [A]t vs. t will yield
a straight line with a slope of -k. Chemical
Kinetics
29.
Determining rxn order
The decomposition of NO2 at 300°C is described by
the equation
NO2 (g) NO (g) + 1/2 O2 (g)
and yields these data:
Time (s) [NO2], M
0.0 0.01000
50.0 0.00787
100.0 0.00649
200.0 0.00481 Chemical
Kinetics
300.0 0.00380
30.
Determining rxn order
Graphing ln [NO2] vs. t yields:
• The plot is not a straight
line, so the process is not
first-order in [A].
Time (s) [NO2], M ln [NO2]
0.0 0.01000 -4.610
50.0 0.00787 -4.845 Does not fit:
100.0 0.00649 -5.038
200.0 0.00481 -5.337 Chemical
Kinetics
300.0 0.00380 -5.573
31.
Second-Order Processes
A graph of 1/[NO2] vs. t
gives this plot.
Time (s) [NO2], M 1/[NO2] • This is a straight
line. Therefore,
0.0 0.01000 100
the process is
50.0 0.00787 127
second-order in
100.0 0.00649 154 [NO2].
200.0 0.00481 208 Chemical
Kinetics
300.0 0.00380 263
32.
• Half-life is defined
as the time required
for one-half of a
reactant to react.
• Because [A] at t1/2 is
one-half of the
original [A],
[A]t = 0.5 [A]0.
Chemical
Kinetics
33.
Half-Life
For a first-order process, set [A]t=0.5 [A]0 in
integrated rate equation:
NOTE: For a first-order
process, the half-life does
Chemical
not depend on [A]0. Kinetics
34.
Half-Life- 2nd order
For a second-order process, set
[A]t=0.5 [A]0 in 2nd order equation.
Chemical
Kinetics
35.
Outline: Kinetics
First order Second order Second order
d Rate complicated
Half-life complicated
Chemical
Kinetics
36.
Temperature and Rate
• Generally, as temperature
increases, so does the
reaction rate.
• This is because k is
temperature dependent.
Chemical
Kinetics
37.
The Collision Model
• In a chemical reaction, bonds are
broken and new bonds are formed.
• Molecules can only react if they collide
with each other.
Chemical
Kinetics
38.
The Collision Model
Furthermore, molecules must collide with the
correct orientation and with enough energy to
cause bond breakage and formation.
Chemical
Kinetics
39.
Activation Energy
• In other words, there is a minimum amount of energy
required for reaction: the activation energy, Ea.
• Just as a ball cannot get over a hill if it does not roll
up the hill with enough energy, a reaction cannot
occur unless the molecules possess sufficient energy
to get over the activation energy barrier.
Chemical
Kinetics
40.
Reaction Coordinate Diagrams
It is helpful to
visualize energy
throughout a
process on a
reaction coordinate
diagram like this
one for the
rearrangement of
methyl isonitrile.
Chemical
Kinetics
41.
Reaction Coordinate Diagrams
• It shows the energy of
the reactants and
products (and,
therefore, E).
• The high point on the
diagram is the transition
state.
• The species present at the transition state is
called the activated complex.
• The energy gap between the reactants and the
activated complex is the activation energy
barrier. Chemical
Kinetics
42.
Maxwell–Boltzmann Distributions
• Temperature is
defined as a
measure of the
average kinetic
energy of the
molecules in a
sample.
• At any temperature there is a wide
distribution of kinetic energies. Chemical
Kinetics
43.
Maxwell–Boltzmann Distributions
• As the temperature
increases, the curve
flattens and
broadens.
• Thus at higher
temperatures, a
larger population of
molecules has
higher energy.
Chemical
Kinetics
44.
Maxwell–Boltzmann Distributions
• If the dotted line represents the activation
energy, as the temperature increases, so does
the fraction of molecules that can overcome
the activation energy barrier.
• As a result, the
reaction rate
increases.
Chemical
Kinetics
45.
Maxwell–Boltzmann Distributions
This fraction of molecules can be found through the expression:
where R is the gas constant and T is the temperature in Kelvin .
Chemical
Kinetics
46.
Arrhenius Equation
Svante Arrhenius developed a mathematical
relationship between k and Ea:
where A is the frequency factor, a number that
represents the likelihood that collisions would
occur with the proper orientation for reaction.
Chemical
Kinetics
47.
Arrhenius Equation
Taking the natural
logarithm of both
sides, the equation
becomes
1
RT
y = mx + b
When k is determined experimentally at
several temperatures, Ea can be calculated
from the slope of a plot of ln k vs. 1/T. Chemical
Kinetics
48.
Outline: Kinetics
First order Second order Second order
d Rate complicated
Half-life complicated
Chemical
Kinetics
49.
Reaction Mechanisms
The sequence of events that describes
the actual process by which reactants
become products is called the reaction
Chemical
Kinetics
50.
Reaction Mechanisms
• Reactions may occur all at once or
through several discrete steps.
• Each of these processes is known as
an elementary reaction or elementary
process.
Chemical
Kinetics
51.
Reaction Mechanisms
• The molecularity of a process tells how many molecules are
involved in the process.
• The rate law for an elementary step is written directly from that step.
Chemical
Kinetics
52.
Multistep Mechanisms
• In a multistep process, one of the steps will
be slower than all others.
• The overall reaction cannot occur faster than
this slowest, rate-determining step.
Chemical
Kinetics
53.
Slow Initial Step
NO2 (g) + CO (g) NO (g) + CO2 (g)
• The rate law for this reaction is found
experimentally to be
Rate = k [NO2]2
• CO is necessary for this reaction to occur, but the
rate of the reaction does not depend on its
concentration.
• This suggests the reaction occurs in two steps.
Chemical
Kinetics
54.
Slow Initial Step
• A proposed mechanism for this reaction is
Step 1: NO2 + NO2 NO3 + NO (slow)
Step 2: NO3 + CO NO2 + CO2 (fast)
• The NO3 intermediate is consumed in the second step.
• As CO is not involved in the slow, rate-determining step, it does
not appear in the rate law.
Chemical
Kinetics
55.
Fast Initial Step
• The rate law for this reaction is found
(experimentally) to be
• Because termolecular (= trimolecular)
processes are rare, this rate law
suggests a two-step mechanism.
Chemical
Kinetics
56.
Fast Initial Step
• A proposed mechanism is
Step 1 is an equilibrium-
it includes the forward and reverse reactions.
Chemical
Kinetics
57.
Fast Initial Step
• The rate of the overall reaction depends
upon the rate of the slow step.
• The rate law for that step would be
• But how can we find [NOBr2]?
Chemical
Kinetics
58.
Fast Initial Step
• NOBr2 can react two ways:
With NO to form NOBr
By decomposition to reform NO and Br2
• The reactants and products of the first
step are in equilibrium with each other.
• Therefore,
Ratef = Rater
Chemical
Kinetics
59.
Fast Initial Step
• Because Ratef = Rater ,
k1 [NO] [Br2] = k−1 [NOBr2]
Solving for [NOBr2] gives us
k1
[NO] [Br2] = [NOBr2]
k−1
Chemical
Kinetics
60.
Fast Initial Step
Substituting this expression for [NOBr2]
in the rate law for the rate-determining
step gives
Chemical
Kinetics
61.
Catalysts
• Catalysts increase the rate of a reaction by
decreasing the activation energy of the
reaction.
• Catalysts change the mechanism by which
the process occurs.
Chemical
Kinetics
62.
Catalysts
One way a
catalyst can
speed up a
reaction is by
holding the
reactants together
and helping bonds
to break.
Chemical
Kinetics
63.
• Enzymes are
catalysts in
biological systems.
• The substrate fits
into the active site of
the enzyme much
like a key fits into a
lock.
Chemical
Kinetics
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