Transcript Fundamentals of General, Organic and Biological Chemistry 6th Edition Chapter Seven Chemical Reactions: Energy, Rates, and Equilibrium James E.

Fundamentals of General,
Organic and Biological
Chemistry
6th Edition
Chapter Seven
Chemical Reactions: Energy,
Rates, and Equilibrium
James E. Mayhugh
Copyright © 2010 Pearson Education, Inc.
Outline
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7.1 Energy and Chemical Bonds
7.2 Heat Changes during Chemical Reactions
7.3 Exothermic and Endothermic Reactions
7.4 Why Do Chemical Reactions Occur? Free Energy
7.5 How Do Chemical Reactions Occur? Reaction Rates
7.6 Effects of Temperature, Concentration, and Catalysts on
Reaction Rates
► 7.7 Reversible Reactions and Chemical Equilibrium
► 7.8 Equilibrium Equations and Equilibrium Constants
► 7.9 Le Châtelier’s Principle: The Effect of Changing
Conditions on Equilibria
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Goals
►1. What energy changes take place during
reactions? Be able to explain the factors that
influence energy changes in chemical reactions.
►2. What is “free energy,” and what is the
criterion for spontaneity in chemistry? Be able to
define enthalpy, entropy, and free-energy changes,
and explain how the values of these quantities affect
chemical reactions.
►3. What determines the rate of a chemical
reaction? Be able to explain activation energy and
other factors that determine reaction rate.
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Goals Contd.
►4. What is chemical equilibrium? Be able to
describe what occurs in a reaction at equilibrium and
write the equilibrium equation for a given reaction.
►5. What is Le Châtelier’s principle? Be able to
state Le Châtelier’s principle and use it to predict the
effect of changes in temperature, pressure, and
concentration on reactions.
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7.1 Energy and Chemical Bonds
► There are two fundamental kinds of energy.
► Potential energy is stored energy. The water in a
reservoir behind a dam, an automobile poised to
coast downhill, and a coiled spring have potential
energy waiting to be released.
► Kinetic energy is the energy of motion. When the
water falls over the dam and turns a turbine, when
the car rolls downhill, or when the spring uncoils and
makes the hands on a clock move, the potential
energy in each is converted to kinetic energy.
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7.2 Heat Changes During Chemical
Reactions
► Bond dissociation energy: The amount of energy
that must be supplied to break a bond and separate
the atoms in an isolated gaseous molecule.
► The triple bond in N2 has a bond dissociation energy
226 kcal/mole, while the single bond in Cl2 has a
bond dissociation energy 58 kcal/mole.
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►Endothermic: A process or reaction that absorbs
heat and has a positive DH.
►Exothermic: A process or reaction that releases
heat and has a negative DH.
►Law of conservation of energy: Energy can be
neither created nor destroyed in any physical or
chemical change.
►Heat of reaction: Represented by DH, is the
difference between the energy absorbed in
breaking bonds and that released in forming bonds.
DH is also known as enthalpy change.
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7.3 Exothermic and Endothermic
Reactions
When the total strength of the bonds formed in the
products is greater than the total strength of the
bonds broken in the reactants, energy is released
and a reaction is exothermic.
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When the total energy of the bonds formed in the
products is less than the total energy of the bonds
broken in the reactants, energy is absorbed and the
reaction is endothermic.
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7.4 Why Do Chemical Reactions
Occur? Free Energy
► Spontaneous process: A process that, once started,
proceeds without any external influence.
► Entropy: The symbol S is used for entropy and it
has the unit of cal/mole·K. The physical state of a
substance and the number of particles have a large
impact on the value of S.
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► Free energy change (DG): Free energy change is used to
describe spontaneity of a process. It takes both DH and DS
into account.
► Exergonic: A spontaneous reaction or process that releases
free energy and has a negative DG.
► Endergonic: A nonspontaneous reaction or process that
absorbs free energy and has a positive DG.
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DG = DH - TDS
DH
DS
DG
(-) favorable
(+) favorable
(-) spontaneous always
(+) unfavorable (-) unfavorable
(+) nonspontaneous always
(-) favorable
(-) spontaneous @ Low T
(+) nonspontaneous @ High T
(-) unfavorable
(+) unfavorable (+) favorable
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(+) nonspontaneous @ Low T
(-) spontaneous @ High T
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7.5 How Do Chemical Reactions
Occur? Reaction Rates
► The value of DG indicates whether a reaction will
occur but it does not say anything about how fast the
reaction will occur or about the details of the
molecular changes that takes place.
► For a chemical reaction to occur, reactant particles
must collide, some chemical bonds have to break,
and new bonds have to form. Not all collisions lead
to products, however.
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One requirement for a productive collision is that the
colliding molecules must approach with the correct
orientation so that the atoms about to form new bonds
can connect.
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Another requirement for a reaction to occur is that the
collision must take place with enough energy to break
the appropriate bonds in the reactant. If the reactant
particles are moving slowly the particles will simply
bounce apart.
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► Activation energy (Ea): The amount of energy
the colliding particles must have for productive
collisions to occur. The size of the activation
energy determines the reaction rate, or how fast
the reaction occurs.
► The lower the activation energy, the greater the
number of productive collisions in a given amount
of time, and faster the reaction.
► The higher the activation energy, the lower the
number of productive collisions, and slower the
reaction.
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7.6 Effects of Temperature, Concentration,
and Catalysts on Reaction Rates
Reaction rates increase with temperature. With more
energy the reactants move faster. The frequency of
collisions and the force with which collisions occur both
increase. As a rule of thumb, a 10°C rise in temperature
causes a reaction rate to double.
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► A second way to speed up a reaction is to increase
the concentrations of the reactants.
► With reactants crowded together, collisions become
more frequent and reactions more likely. Flammable
materials burn more rapidly in pure oxygen than in
air because the concentration of molecules is higher
(air is approximately 21% oxygen).
► Hospitals must therefore take extraordinary
precautions to ensure that no flames are used near
patients receiving oxygen.
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► A third way to speed up a reaction is to add a
catalyst—a substance that accelerates a chemical
reaction but is itself unchanged in the process.
► A catalyzed reaction has a lower activation energy.
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The thousands of biochemical reactions continually
taking place in our bodies are catalyzed by large protein
molecules called enzymes, which promote reaction by
controlling the orientation of the reacting molecules.
Since almost every reaction is catalyzed by its own
specific enzyme, the study of enzyme structure, activity,
and control is a central part of biochemistry.
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7.7 Reversible Reactions and Chemical
Equilibrium
Imagine the situation if you mix acetic acid and ethyl
alcohol. The two begin to form ethyl acetate and
water. But as soon as ethyl acetate and water form,
they begin to go back to acetic acid and ethyl
alcohol. Such a reaction, which easily goes in either
direction, is said to be reversible and is indicated by
a double arrow in equations.
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Both reactions occur until the concentrations of reactants
and products reach constant values. The reaction vessel
contains both reactants and products and is said to be in a
state of chemical equilibrium. A state in which the rates of
forward and reverse reactions are the same.
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7.8 Equilibrium Equations and
Equilibrium Constants
► Consider the following general equilibrium reaction:
aA + bB + …  mM + nN + …
► Where A, B, … are the reactants; M, N, …. Are the
products; a, b, ….m, n, …. are coefficients in the
balanced equation. At equilibrium, the composition
of the reaction mixture obeys an equilibrium
equation.
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► The equilibrium constant K is the number obtained
by multiplying the equilibrium concentrations of the
products and dividing by the equilibrium
concentrations of the reactants, with the
concentration each substance raised to a power equal
to its coefficient in the balanced equation.
► The value of K varies with temperature.
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► K larger than 1000: Reaction goes essentially to
completion.
► K between 1 and 1000: More products than
reactants are present at equilibrium.
► K between 1 and 0.001: More reactants than
products are present at equilibrium.
► K smaller than 0.001: Essentially no reaction
occurs.
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7.9 Le Châtelier's Principle: The Effect
of Changing Conditions on Equilibria
► Le Châtelier's Principle: When a stress is applied
to a system at equilibrium, the equilibrium shifts to
relieve the stress.
► The stress can be any change in concentration,
pressure, volume, or temperature that disturbs
original equilibrium.
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► What happens if the concentration of CO is
increased?
► To relieve the “stress” of added CO, according to Le
Châtelier’s principle, the extra CO must be used up.
In other words, the rate of the forward reaction must
increase to consume CO.
► Think of the CO added on the left as “pushing” the
equilibrium to the right:
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► The forward and reverse reaction rates adjust until they
are again equal and equilibrium is reestablished.
► At this new equilibrium state, the value of [H2] will be
lower, because more has reacted with the added CO,
and the value of [CH3OH] will be higher.
► The changes offset each other, however, so the value of
the equilibrium constant K remains constant.
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► Le Châtelier’s principle predicts that an increase in
temperature will cause an equilibrium to shift in
favor of the endothermic reaction so the additional
heat is absorbed.
► You can think of heat as a reactant or product whose
increase or decrease stresses an equilibrium just as a
change in reactant or product concentration does.
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► Pressure influences an equilibrium only if one or
more of the substances involved is a gas. As
predicted by Le Châtelier’s principle, increasing the
pressure shifts the equilibrium in the direction that
decreases the number of molecules in the gas phase
and thus decreases the pressure.
► For the ammonia synthesis, increasing the pressure
favors the forward reaction because 4 moles of gas is
converted to 2 moles of gas.
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The effects of changing reaction conditions on
equilibria are summarized below.
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Chapter Summary
► The strength of a covalent bond is measured by its
bond dissociation energy.
► If heat is released, DH is negative and the reaction is
said to be exothermic. If heat is absorbed, DH is
positive and the reaction is said to be endothermic.
► Spontaneous reactions are those that, once started,
continue without external influence; nonspontaneous
reactions require a continuous external influence.
► Spontaneity depends on two factors, the amount of
heat absorbed or released in a reaction and the
entropy change.
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Chapter Summary Contd.
► Spontaneous reactions are favored by a release of heat,
DH <0, and an increase in entropy, DS >0.
► The free-energy change, DG = DH - T DS, takes both
factors into account.
► DG<0 indicates spontaneity, DG>0 indicates
nonspontaneity.
► Chemical reactions occur when reactant particles collide
with proper orientation and energy. The exact amount of
collision energy necessary is the activation energy.
► Reaction rates can be increased by raising the
temperature, by raising the concentrations of reactants,
or by adding a catalyst.
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Chapter Summary Contd.
► At equilibrium, the forward and reverse reactions occur
at the same rate, and the concentrations of reactants
and products are constant. Every reversible reaction
has an equilibrium constant, K. The forward reaction is
favored if K>1; the reverse reaction is favored if K<1.
► Le Châtelier’s principle states that when a stress is
applied to a system in equilibrium, the equilibrium
shifts so that the stress is relieved.
► Applying this principle allows prediction of the effects
of changes in temperature, pressure, and concentration.
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Key Words
►Activation energy
►Bond dissociation
energy
►Catalyst
►Chemical equilibrium
►Endergonic
►Endothermic
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►Enthalpy
►Enthalpy change
►Entropy (S)
►Equilibrium constant
(K)
►Exergonic
►Exothermic
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Key Words Contd.
►Free-energy change
►Heat
►Heat of reaction
►Kinetic energy
►Law of conservation of
energy
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►Le Châtelier’s
principle
►Potential energy
►Reaction rate
►Reversible reaction
►Spontaneous process
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End of Chapter 7
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