Download how reactions occur

SPONTANEOUS PROCESSES
• Spontaneous processes are processes that take place
naturally with no apparent cause or stimulus.
EXERGONIC PROCESSES
• Exergonic processes are processes that give up energy
as they take place. Energy appears on the right side of
equations representing the processes.
ENDERGONIC PROCESSES
• Endergonic processes are processes that gain or absorb
energy as they take place. Energy appears on the left
side of equations representing the processes.
ENTROPY
• Entropy is a measurement or indication of the disorder or
randomness of a system. The more disorderly or mixed up
a system is, the higher its entropy.
FACTORS THAT INFLUENCE SPONTANEITY
• A process will always be spontaneous if it gives up energy
and the entropy of the system involved in the process
increases.
• A process that absorbs energy will be spontaneous only if
an increase in entropy of the system also occurs that is
large enough to compensate for the increase in energy.
• A process that causes an entropy decrease in the system
will be spontaneous only if a decrease in energy also
occurs that is large enough to compensate for the entropy
decrease.
STABLE SUBSTANCES
• Stable substances are substances that do not undergo
spontaneous changes at the prevailing conditions.
REACTION RATES
• A reaction rate is the speed of a reaction.
• Reaction rates can be determined experimentally by
measuring the change in concentration of a reactant or
product and dividing the change by the time required for
the change to occur, using the following equation.
• In this equation, ∆C is the change in concentration of a
reactant or a product that occurs in a measured amount of
time, ∆t. The value of ∆C is calculated by subtracting the
initial concentration, C0, from the final concentration, Ct.
CALCULATION EXAMPLE
• Calculate the average rate of the following reaction:
Na(s) + H2O(l)
H2(g) + NaOH(aq)
A piece of solid Na was put into pure water. All the Na had
reacted after after 8 seconds, and the concentration of NaOH
was found to be 0.0013 M.
• Solution: The initial concentration of NaOH in the water,
C0, was zero because no reaction had occurred before the
Na was added. The change in concentration, ∆C, is equal
to Ct-C0 or 0.0013M -0. This equals 0.0013M. The time of
the reaction was 8 seconds. The rate is calculated as
follows:
HOW REACTIONS OCCUR
• An explanation of how reactions occur is called a reaction
mechanism. A reaction mechanism is often expressed as
a number of processes that must take place for reactions
to occur. The following assumptions make up a reaction
mechanism.
• Assumption 1: Reactant particles must collide with one
another in order for a reaction to occur.
• Assumption 2: Reactant particles must collide with at least
a certain total energy if the collision is to result in a
reaction.
• In some cases, colliding reactant particles must be
oriented in a specific way if a reaction is to occur.
WHY COLLISIONS ARE NECESSARY
• Reactant particles must collide if they are to react
(assumption 1). With few exceptions, molecules cannot
react with each other if they do not come in contact.
During collisions, some bonds are broken, atoms are
exchanged and new bonds form.
WHY MINIMUM COLLISION ENERGY IS NECESSARY
• The requirement that some bonds of reactant molecules
must break if a reaction is to occur makes the requirement
of minimum collision energies valid.
• At all temperatures above absolute zero, the bonded
atoms of molecules are vibrating and stretching the bonds.
The energy associated with these vibrations is called
internal energy. Internal energy can be increased by
collisions. This results in an increase in the vibrational
amplitude which increases the chances for a bond to
break.
• These ideas are illustrated in the following drawing:
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.
• 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.
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.
• 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. This idea is illustrated by
the following drawing:
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.
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.
DIAGRAMS SHOWING DIFFERENCES IN ACTIVATION
ENERGY
• Activation energy differences become quite obvious in
energy diagrams as shown by the following illustrations:
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, and the presence of catalysts.
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.
• 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. For example, if molecules A and B are
to react, a collision between an A molecule and a B
molecule could not possibly lead to 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.
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.
• 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 finelydivided iron powder rusts much more rapidly because the
surface area and effective concentration is much greater.
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.
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.
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.
• 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.
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.
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.
EQUILIBRIUM EXPRESSION
• The equilibrium expression for the above equilibrium 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.
• This is demonstrated for the following equilibrium:
2NO(g) + 2H2(g)
N2(g) + 2H2O(g)
EQUILIBRIUM CONSTANT
• The K in both 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 indicates an equilibrium position
toward the left or reactant side of the equilibrium.
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.
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.
• 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.
• 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.