Transcript Reaction Rates

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Table of Contents
Topic
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Topic 21: Reaction Rates
Basic Concepts
Additional Concepts
Reaction Rates: Basic Concepts
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Exothermic and Endothermic Reactions
• An exothermic reaction releases heat, and an
endothermic reaction absorbs heat.
• The exothermic
reaction gives off
heat because the
products are at a
lower energy level
than the reactants.
Reaction Rates: Basic Concepts
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Exothermic and Endothermic Reactions
• The endothermic reaction absorbs heat
because the
products are at
a higher energy
level than the
reactants.
Reaction Rates: Basic Concepts
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Exothermic and Endothermic Reactions
• Scientists have observed that the energy
released in the formation of a compound from
its elements is always identical to the energy
required to decompose that compound into its
elements.
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Exothermic and Endothermic Reactions
• This observation is an illustration of an
important scientific principle known as the
law of conservation of energy. That law
states that energy is neither created nor
destroyed in a chemical change, but is simply
changed from one form to another.
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Exothermic and Endothermic Reactions
• In an exothermic reaction, the heat released
comes from the change from reactants at
higher energy to products at lower energy.
• In an endothermic reaction, the heat absorbed
comes from the opposite change.
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Expressing Reaction Rates
• As you know, some chemical reactions are
fast and others are slow; however, fast and
slow are inexact, relative terms.
• Chemists, engineers, medical researchers,
and others often need to be more specific.
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Expressing Reaction Rates
• We generally define the average rate of an
action or process to be the change in a given
quantity during a specific period of time.
• Recall from your study of math that the
Greek letter delta (∆) before a quantity
indicates a change in the quantity.
• In equation form, average rate or speed is
written as
Reaction Rates: Basic Concepts
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Expressing Reaction Rates
• For chemical reactions, this equation defines
the average rate at which reactants produce
products, which is the amount of change of a
reactant in a given period of time.
• Most often, chemists are concerned with
changes in the molar concentration (mol/L,
M) of a reactant or product during a reaction.
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Expressing Reaction Rates
• Therefore, the reaction rate of a
chemical reaction is stated as the change
in concentration of a reactant or product
per unit time, expressed as mol/(L∙s).
• Brackets around the formula for a substance
denote the molar concentration. For example,
[NO2] represents the molar concentration of
NO2.
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Expressing Reaction Rates
• It’s important to understand that reaction rates
are determined experimentally by measuring
the concentrations of reactants and/or
products in an actual chemical reaction.
• Reaction rates cannot be calculated from
balanced equations as stoichiometric
amounts can.
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Expressing Reaction Rates
• Reaction rates must always be positive.
• When the rate is measured by the
consumption of a reactant, scientists
commonly apply a negative sign to the
calculation to get a positive reaction rate.
• Therefore, the following form of the average
rate equation is used to calculate the rate of
consumption:
Reaction Rates: Basic Concepts
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Calculating Average Reaction Rates
• Reaction data for the reaction between butyl
chloride (C4H9Cl) and water (H2O) is given in
the table.
• Calculate the average reaction rate over this
time period
expressed as
moles of C4H9Cl
consumed per
liter per second.
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Calculating Average Reaction Rates
• You are given the initial and final
concentrations of C4H9Cl and the initial and
final times.
• You can calculate the average reaction rate
of the chemical reaction using the change in
concentration of butyl chloride in four
seconds.
• In this problem, the reactant butyl chloride is
consumed.
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Calculating Average Reaction Rates
• Write the equation for the average reaction
rate, insert the known quantities, and
perform the calculation.
Reaction Rates: Basic Concepts
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Calculating Average Reaction Rates
Reaction Rates: Basic Concepts
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Collision Theory
• According to the collision theory, atoms,
ions, and molecules must collide with each
other in order to react.
• The following three statements summarize
the collision theory.
1. Particles must collide in order to react.
2. The particles must collide with the correct
orientation.
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Collision Theory
3. The particles must collide with enough
energy to form an unstable activated
complex, also called a transition state,
which is an intermediate particle made up
of the joined reactants.
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Collision Theory
• The minimum amount of energy that
colliding particles must have in order to form
an activated complex is called the activation
energy of the reaction.
• Particles that
collide with less
than the activation
energy cannot
form an activated
complex.
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Reaction Rates: Basic Concepts
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Collision Theory
• In an exothermic reaction, molecules collide
with enough energy to overcome the
activation energy barrier, form an activated
complex, then
release energy and
form products at a
lower energy
level.
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Collision Theory
• In the reverse endothermic reaction, the
reactant molecules lying at a low energy level
must absorb
energy to
overcome the
activation
energy barrier
and form highenergy
products.
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Collision Theory
• Reaction spontaneity is related to change
in free energy, ∆G.
• If ∆G is negative, the reaction is
spontaneous under the conditions specified.
• If ∆G is positive, the reaction is not
spontaneous.
• Let’s now consider whether spontaneity has
any effect on reaction rates.
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Collision Theory
• To investigate the relationship between
spontaneity and reaction rate, consider the
following gas-phase reaction between
hydrogen and oxygen.
• Here, ∆G = –458 kJ at 298 K (25°C) and 1
atm pressure.
• Because ∆G is negative, the reaction is
spontaneous.
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Collision Theory
• For the same reaction, ∆G = 2484 kJ, which
means that the reaction is highly exothermic.
• You can examine the
speed of this reaction by
filling a tape-wrapped
soda bottle with
stoichiometric quantities
of the two gases—two
volumes hydrogen and one
volume oxygen.
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Collision Theory
• A thermometer in the stopper allows you to
monitor the temperature inside the bottle.
• As you watch for
evidence of a reaction,
the temperature remains
constant for hours.
• Have the gases
escaped? Or have they
simply failed to react?
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Collision Theory
• If you remove the stopper and hold a burning
splint to the mouth of the bottle, a reaction
occurs explosively.
• Clearly, the hydrogen and
oxygen gases have not
escaped from the bottle.
• Yet they did not react
noticeably until you supplied
additional energy in the form
of a lighted splint.
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Collision Theory
• As these examples show, reaction spontaneity
in the form of ∆G implies absolutely nothing
about the speed of the reaction; ∆G merely
indicates the natural tendency for a reaction
or process to proceed.
• Factors other than spontaneity, however, do
affect the rate of a chemical reaction.
• These factors are keys in controlling and
using the power of chemistry.
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Factors Affecting Reaction Rates
• The reaction rate for almost any chemical
reaction can be modified by varying the
conditions of the reaction.
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The Nature of Reactants
• An important factor that affects the rate of a
chemical reaction is the reactive nature of the
reactants. As you know, some substances
react more readily than others.
• The tendency of a substance to react
influences the rate of a reaction involving
the substance.
• The more reactive a substance is, the faster
the reaction rate.
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Concentration
• Reactions speed up when the concentrations
of reacting particles are increased.
• One of the fundamental principles of the
collision theory is that particles must collide
to react.
• The number of particles in a reaction makes
a difference in the rate at which the reaction
takes place.
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Concentration
• Think about a reaction where reactant A
combines with reactant B.
• At a given concentration of A and B, the
molecules of A collide with B to produce
AB at a particular rate.
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Concentration
• What happens if the amount of B is increased?
• Increasing the concentration of B makes more
molecules available with which A can collide.
• Reactant A “finds” reactant B more easily
because there are more B molecules in the
area, which increases probability of collision,
and ultimately increases the rate of reaction
between A and B.
Reaction Rates: Basic Concepts
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Surface Area
• Increasing the surface area of reactants
provides more opportunity for collisions with
other reactants,
thereby
increasing the
reaction rate.
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Temperature
• Generally, increasing the temperature at which
a reaction occurs increases the reaction rate.
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Temperature
• As you can see, increasing the temperature of
the reactants increases the reaction rate
because raising the
kinetic energy of
the reacting
particles raises
both the collision
frequency and the
collision energy.
Reaction Rates: Basic Concepts
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Temperature
• When the temperature is increased, more of
the collisions at high temperatures have
enough energy to
overcome the
activation energy
barrier and react.
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Catalysts
• You’ve seen that increasing the temperature
and/or the concentration of reactants can
dramatically increase the rate of a reaction.
• However, an increase in temperature is
not always the best (or most practical)
thing to do.
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Catalysts
• For example, suppose that you want to
increase the rate of a reaction such as the
decomposition of glucose in a living cell.
• Increasing the temperature and/or the
concentration of reactants is not a viable
alternative because it might harm or kill
the cell.
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Catalysts
• It is a fact that many chemical reactions in
living organisms would not occur quickly
enough to sustain life at normal living
temperatures if it were not for the presence
of enzymes.
• An enzyme is a type of catalyst, a substance
that increases the rate of a chemical reaction
without itself being consumed in the
reaction.
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Catalysts
• Although catalysts are important substances
in a chemical reaction, a catalyst does not
yield more product and is not included in the
product(s) of the reaction.
• In fact, catalysts are not included in the
chemical equation.
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Catalysts
• This energy diagram shows how the activation
energy of the catalyzed reaction is lower and
therefore the
reaction
produces the
products at a
faster rate
than the
uncatalyzed
reaction does.
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Catalysts
• Note that the activation energy for the
catalyzed reaction is much lower than for the
uncatalyzed reaction.
• The lower activation energy for the catalyzed
reaction means that more collisions have
sufficient energy to initiate reaction.
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Catalysts
• Another type of substance that affects reaction
rates is called an inhibitor.
• Unlike a catalyst, which speeds up reaction
rates, an inhibitor is a substance that slows
down, or inhibits, reaction rates.
• Some inhibitors, in fact, actually prevent a
reaction from happening at all.
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Heterogeneous and
homogeneous catalysts
• A heterogeneous catalyst exists in a physical
state different than that of the reaction it
catalyzes.
• Because the catalysts in a catalytic converter
are solids and the reactions they catalyze are
gaseous, the catalysts are called heterogeneous
catalysts.
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Heterogeneous and
homogeneous catalysts
• A catalyst that exists in the same physical
state as the reaction it catalyzes is called a
homogeneous catalyst.
• For example, if both an enzyme and the
reaction it catalyzes are in aqueous solution,
the enzyme is a homogeneous catalyst.
Basic Assessment Questions
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Question 1
In aqueous solution, bromine reacts with
formic acid (HCOOH) to produce hydrogen
bromide and carbon dioxide gas. Calculate
the average reaction rate over 150 s expressed
in terms of the disappearance of Br2.
Basic Assessment Questions
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Answer
Basic Assessment Questions
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Question 2
Use the factors discussed in this section to
account for the following observation.
When two gases react, compressing the gases
generally increases the rate of reaction.
Basic Assessment Questions
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Answer
Basic Assessment Questions
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Question 3
Use the factors discussed in this section to
account for the following observation.
The rate of gaseous reactions can also be
increased by pumping more gas into the
reaction container.
Basic Assessment Questions
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Answer
Adding gas increases concentration by adding
particles to a given space.
Reaction Rates: Additional Concepts
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Additional Concepts
Reaction Rates: Additional Concepts
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Reaction Rate Laws
• The equation that expresses the mathematical
relationship between the rate of a chemical
reaction and the concentration of reactants is
called a rate law.
• For example, the reaction A → B, which is a
one-step reaction, has only one activated
complex between reactants and products.
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Reaction Rate Laws
• The rate law for this reaction is expressed as,
Rate = k [A] where k is the specific rate
constant, or a numerical value that relates
reaction rate and concentration of reactants at
a given temperature.
Reaction Rates: Additional Concepts
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•
•
•
•
Reaction Rate Laws
Units for the rate constant include L/(mol∙s),
L2/(mol2∙s), and s–1.
Depending on the reaction conditions,
especially temperature, k is unique for
every reaction.
The rate law means that the reaction rate is
directly proportional to the molar
concentration of A.
Thus, doubling the concentration of A will
double the reaction rate.
Reaction Rates: Additional Concepts
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•
•
•
•
Reaction Rate Laws
Increasing the concentration of A by a factor
of 5 will increase the reaction rate by a
factor of 5.
The specific rate constant, k, does not change
with concentration; however, k does change
with temperature.
A large value of k means that A reacts rapidly
to form B.
What does a small value of k mean?
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Reaction order
• In the expression Rate = k[A], it is
understood that the notation [A] means the
same as [A]1.
• In other words, for reactant A, the understood
exponent 1 is called the reaction order.
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Reaction order
• The reaction order for a reactant defines
how the rate is affected by the concentration
of that reactant.
• For example, the rate law for the
decomposition of H2O2 is expressed by the
following equation.
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Reaction order
• Because the reaction rate is directly
proportional to the concentration of H2O2
raised to the first power, [H2O2]1, the
decomposition of H2O2 is said to be first
order in H2O2.
• Because the reaction is first order in H2O2,
the reaction rate changes in the same
proportion that the concentration of H2O2
changes.
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Reaction order
• So if the H2O2 concentration decreases to
one-half its original value, the reaction rate
is halved as well.
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Reaction order
• Because reaction order is based on reaction
rates, it follows that reaction order also is
determined experimentally.
• Finally because the rate constant describes
the reaction rate, k, too, must be determined
experimentally.
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Determining Rate Laws
• Consider the following reaction and its
experimentally determined rate law.
• This rate equation shows that the rate
depends on the concentrations of both NO2
and F2, each to the first power.
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Determining Rate Laws
• In other words, the reaction is first order in
NO2 and first order in F2.
• So, if [NO2] is doubled while [F2] remains
the same, the rate doubles.
• Also, if [F2] is doubled while [NO2] remains
the same, the reaction rate doubles.
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Determining Rate Laws
• Now, examine the following reaction and its
rate law.
• Because the rate depends on the square of the
concentration of NO, doubling [NO] while
leaving [O2] the same will increase the
reaction rate by a factor of 22, or 4.
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Determining Rate Laws
• Use the data in the table below to determine
the form of the rate law for the following
reaction.
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Determining Rate Laws
• The general rate law for this type of reaction
is as follows.
• To start compare the data from trial 1 and 2.
• Notice that [NO] in trial 1 is 4.0 x 10–3 ,
mo1/L and [NO] in trial 2 is 8.0 x 10–3
mo1/L, double that in trial 1.
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Determining Rate Laws
• Now see how the rate changes from trial 1 to
trial 2.
• The rate in trial 2, 4.8 x 10–5 mo1/L ∙s, is four
times the rate in trial 1, which is 1.2 x 10–5
mo1/L∙s.
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Determining Rate Laws
• When [NO] doubles, the initial reaction rate
quadruples.
• Therefore, it is likely that the reaction rate
depends on the square of the concentration
of NO.
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Determining Rate Laws
• Next, determine how the rate depends on the
change in [H2].
• When [H2] doubles, the initial rate doubles.
• This result indicates that the rate is directly
proportional to the concentration of H2.
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Determining Rate Laws
• Now you can write the rate law based on
your comparisons.
• The rate law means that the reaction is
second order in [NO] and first order in [H2].
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Instantaneous Reaction Rates
and Reaction Mechanisms
• Many chemical reactions are complex.
• In a complex reaction, a certain series
of steps must occur with correct fit and in
an exact sequence in order to yield the
reaction products. This series of steps is
called a reaction mechanism.
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Instantaneous Reaction Rates
and Reaction Mechanisms
• As an example, consider that hydrogen
peroxide (H2O2) decomposes when iodide
ions (I–) are present.
• Experiments have shown that the reaction
takes place in two steps as follows.
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Instantaneous Reaction Rates
and Reaction Mechanisms
• The IO– ion is called an intermediate in the
reaction.
• An intermediate is an atom, an ion, or a
molecule produced in one step of a reaction
and consumed in a later step.
Reaction Rates: Additional Concepts
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Instantaneous Reaction Rates
and Reaction Mechanisms
• In a complex reaction, one step is always
slower than the others.
• This step is called the rate-determining
step because it will determine how fast the
reaction forms products, no matter how fast
the other steps are.
• In the reaction, the first step is the slower
step, therefore, it is the rate-determining step
for the reaction
Additional Assessment Questions
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Question 1
The rate law for a reaction
is
rate = k[A][B]. What happens to the rate if
the concentration of B is tripled while [A]
remains the same? If the concentration of A
is halved while [B] remains the same?
Additional Assessment Questions
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Answer
rate triples; rate halves
Additional Assessment Questions
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Question 2
The rate law for a reaction
is
rate = k[X][Y]2. What happens to the rate if
the concentration of X is doubled while [Y]
remains the same?
Additional Assessment Questions
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Answer
rate doubles
Additional Assessment Questions
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Question 3
Discuss two possible ways to determine the
average rate of the following reaction in the
laboratory. What quantities would you
measure in each case?
Additional Assessment Questions
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Answer
There are three possibilities. Measure the
reduction in concentration of H2, the reduction
in concentration of I2, or the increase in
concentration of HI. In all three cases, time is
the second quantity that must be measured.
Additional Assessment Questions
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Question 4
The collision theory says that particles must
collide in order to react. What other two
requirements must be met in order for particles
to react?
Additional Assessment Questions
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Answer
The particles must collide in the correct
orientation and with sufficient energy to form
an activated complex.
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