Chapter 8:
Reaction Rates and Equilibrium
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ACTIVATION ENERGY
• In some reaction mixtures, the average total energy of the
molecules is too low at the prevailing temperature for a
reaction to take place at a detectable rate. The reaction
mixture is said to be stable.
• For many stable mixtures, the addition of a small amount
of energy starts the reaction which then continues without
the addition of any more stimulus or energy from an
outside source.
• The small amount of outside
energy needed to start
spontaneous processes is
called activation energy.
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• An ordinary kitchen match provides a good example of
these concepts. The reactants in the match head are
stable until the match is rubbed on a rough surface. The
energy of rubbing provides the necessary activation
energy to cause the match head components to react and
the match ignites. Once ignited, the match continues to
burn spontaneously until all the fuel of the match has
reacted with oxygen in the air.
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MOLECULAR ORIENTATION
• Orientation effects are related to which side or end of a
reacting particle actually contacts another particle during a
collision.
• The orientation of reacting particles is not important in
some reactions such as those between reacting ions in a
solution. For example, Ca2+ ions react with CO32- in
solution to form solid CaCO3 is insoluble and settles out of
the solution. Both ions can be considered
to be spherical charged particles, so their
orientation toward each other when they
collide does not influence the reaction
rate.
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• Orientation effects become important when the reacting
particles are not spherical as in the following hypothetical
reaction:
A-B + C-D
A-C + B-D
• In this reaction it is obvious that
the chances for a reaction to occur
are better if the A-end of the first
molecule hits the C-end of the
second molecule.
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ENERGY DIAGRAMS
• Energy relationships for reactions can be illustrated
graphically by energy diagrams in which the energy of a
reaction is graphed on the vertical axis versus the progress
of the reaction on the horizontal axis. A general energy
diagram is shown below with important quantities labeled.
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EXOTHERMIC AND ENDOTHERMIC DIAGRAMS
• The difference between endothermic and exothermic
reactions is clearly indicated by the following energy
diagrams. Note that in exothermic reactions the energy is
lost as the reaction occurs. Hence, the products have less
energy than the reactants. The reverse is true for
endothermic reactions which gain energy and cause the
products to have more energy than reactants.
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DIAGRAMS SHOWING DIFFERENCES IN ACTIVATION
ENERGY
• Activation energy differences become quite obvious in
energy diagrams as shown by the following illustrations:
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FACTORS THAT INFLUENCE REACTION RATES
• Reaction rates are influenced by a number of different
factors, including:
• the nature of the reactants,
• the concentration of the reactants,
• the temperature of the reactants,
• the presence of catalysts.
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THE NATURE OF THE REACTANTS
• Reactions between oppositely-charged ions in solution
occur almost instantaneously. This is because the ions are
strongly attracted to each other because of their opposite
electrical charges.
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• Reactions between covalently-bonded molecules in which
covalent bonds have to be broken often take place slowly.
This is partially because the molecules collide only
because of their random motion, and all collisions do not
result in a reaction.
• Other characteristics of reactants such as their physical
state (gases, liquids or solids), their molecular sizes, and
whether or not they are polar are also important influences
of some reaction rates.
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THE CONCENTRATION OF THE REACTANTS
• The requirement for a collision to occur between reactant
molecules before a reaction can take place accounts for
the reactant concentration influence on reaction rates.
• If a reaction occurs between A and B molecules, and a
reaction mixture contains mostly A molecules, most
collisions participated in by A molecules will be with other A
molecules and the reaction rate will be low.
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• The reaction between a solid piece of iron and oxygen gas
takes place slowly in part because only iron atoms on the
surface can collide with oxygen molecules. The effective
concentration of iron is low. However finely-divided iron
powder rusts much more rapidly because the surface area
and effective concentration is much greater.
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THE TEMPERATURE OF THE REACTANTS
• The effect of temperature on reaction rates can also be
explained using the concept of molecular collisions.
• An increase in the temperature of the reactants
corresponds to an increase in the velocity and the kinetic
energy of the molecules.
• An increase in velocity will increase the number of
molecular collisions that take place in a fixed amount of
time and will thus increase the rate of the reaction.
• An increase in the kinetic energy of the colliding molecules
will increase the internal energy of the molecules and also
increase the number of molecules with the required
minimum activation energy.
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THE PRESENCE OF CATALYSTS
• Catalysts are substances that speed up chemical reactions
without being used up in the reaction.
• Homogeneous catalysts are substances that are
distributed uniformly throughout a reaction mixture.
• Heterogeneous catalysts are substances normally used in
the form of solids with large surface areas on which the
reactions take place.
• One explanation for catalytic behavior is that catalysts
provide an alternate reaction pathway that requires less
activation energy than the normal pathway.
• Another explanation proposes that solid catalysts provide a
surface on which reactant molecules adsorb with favorable
orientations to each other. Adsorbed molecules with
favorable orientations are located close enough to each
other to react rapidly.
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CHEMICAL EQUILIBRIUM
• All chemical reactions can (in principle) go in both
directions and products, located to the right of the arrow,
can react to form reactants, located to the left of the arrow.
This condition is indicated by the use of a double arrow
pointing in both directions as shown below:
H2(g) + I2(g)
2HI(g)
• When the rate of the reaction toward the right is equal to
the rate of the reaction toward the left, the reaction is said
to be in equilibrium.
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• When a reaction is in equilibrium, the concentrations of
reactants and products remain constant as time passes.
• The unchanging concentrations of reactants and products
in a reaction at equilibrium are called equilibrium
concentrations.
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THE POSITION OF EQUILIBRIUM
• The position of equilibrium is an indication of the relative
amounts of reactants and products present in a reaction
mixture at equilibrium.
• The position is said to be to the right when the amount of
product is significantly more than the amount of reactant.
• The position is to the left when more reactant is present
than product.
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MATHEMATICAL REPRESENTATION OF POSITION
OF EQUILIBRIUM
• The position of equilibrium can be represented
mathematically by using the concepts of an equilibrium
expression and an equilibrium constant.
• Both concepts will be initially represented using the
following hypothetical equilibrium:
aA + bB
wW + xX
• In this expression, the lower case letters represent the
stoichiometric coefficients of the reaction and the upper case
letters represent the formulas of the reacting substances.
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EQUILIBRIUM EXPRESSION
• The equilibrium expression for the general equation on the
previous slide is written as follows:
• In this equation, the brackets,[ ], stand for molar
concentrations of the reactants, A and B, and the products, W
and X. It is seen that each reactant concentration is raised to
a power equal to the stoichiometric coefficient of that reactant
in the equilibrium equation.
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• This is demonstrated for the following equilibrium:
2NO(g) + 2H2(g)
N2(g) + 2H2O(g)
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EQUILIBRIUM CONSTANT
• The K in equilibrium expressions is called the equilibrium
constant.
• As long as the temperature does not change, it has a
constant value because none of the concentrations used to
express it change with time once equilibrium is
established.
• A relatively large value for K indicates that the equilibrium
position is toward the right or products side of the
equilibrium.
• A small K indicates an equilibrium position toward the left
or reactant side of the equilibrium.
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THE RANGE OF K VALUES
• The values for K that have been found experimentally
range between wide extremes.
• Some, such as, K= 1.1 x 10-36, are so small that for all
practical purposes an equilibrium mixture would contain only
reactants and the equilibrium position is extremely far to the
left.
• Others, such as K= 1.2 x1040, are so large that for all practical
purposes an equilibrium mixture would contain only products
and the equilibrium position is extremely far to the right.
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FACTORS THAT INFLUENCE THE POSITION OF
EQUILIBRIUM
• According to Le Châtelier's principle the position of an
equilibrium shifts in response to changes made in the
equilibrium.
• The factors that will be considered are:
•concentrations of reactants and products
•reaction temperature
•catalysts
• In general, Le Châtelier's principle predicts a shift away from
the side to which something is added and toward the side
from which something is removed.
• Catalysts cannot change the position of an equilibrium
because it lowers the energy barrier for both the forward and
reverse reactions; therefore, catalysts speed up both
forward and reverse reactions and cannot change the
position of equilibrium.
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• Consider the following endothermic reaction at equilibrium:
heat + 4NO2(g) + 6H2O(g)
7O2 (g) + 4NH3(g)
• If an equilibrium mixture was heated, the equilibrium
position would shift toward the right to try to use up the
added heat.
• If some NO2 was added to an equilibrium mixture, the
equilibrium position would again shift toward the right to try
to use up the added NO2.
• If some NH3 was removed from an equilibrium mixture, the
equilibrium position would again shift toward the right in an
attempt to replace the NH3 that was removed.
• If a catalyst were added, the equilibrium position would not
change.
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