Spencer L. Seager
Michael R. Slabaugh www.cengage.com/chemistry/seager
Jennifer P. Harris

• 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.
2HgO (s) + 4.34 kcal → 2Hg (l) + O
2
(g)

• 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 .

• 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 . (e.g. Ice spontaneously melts at 20 °C.)
• 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. (e.g. Water spontaneously freezes at 0 °C.)

• Stable substances are substances that do not undergo spontaneous changes at the prevailing conditions.
The three states of matter for bromine ( a ) bromine as a solid, ( b ) bromine as a liquid, ( c ) bromine as a gas.

• 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:
Rate



C t

C t

Δt
C
0
• 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, C
0
, from the final concentration, C t
.

• Calculate the average rate of the following reaction:
2 Na(s) + 2 H
2
O(l) H
2
(g) + 2 NaOH(aq)
A piece of solid Na was put into pure water. All the Na had reacted after 8 seconds, and the concentration of NaOH was found to be 0.0013 M.
• Solution: The initial concentration of NaOH in the water, C
0
, was zero because no reaction had occurred before the Na was added. The change in concentration, ∆C, is equal to C t
-
C
0 or 0.0013M - 0. This equals 0.0013M. The time of the reaction was 8 seconds. The rate is calculated as follows:
Rate


C
 t

1.3

10
-3
M
8 sec

2

10

4
M sec

• 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.
• Assumption 3: 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:

• 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.

• 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, Ca 2+ ions react with CO
3
2in solution to form solid
CaCO
3
, which 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.

• 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.

• 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.

• Activation energy differences become quite obvious in energy diagrams as shown by the following illustrations:

• 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 or inhibitors .

• 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.
• 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 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 CONCENTRATION OF THE REACTANTS
(continued)
• 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, wire wool and finelydivided iron powder rust much more rapidly because the surface area and effective concentration is much greater.

• An effective collision is a collision that causes a reaction to occur between the colliding molecules.
• An effective collision is a collision:
• between the reactants for the desired chemical reaction.
• with the right amount of energy.
• with the right orientation.

• 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 reaction rate .
• 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 .

DISTRIBUTION OF KINETIC ENERGIES
(MOLECULAR VELOCITIES) AT TWO TEMPERATURES

• 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.

• An inhibitor is a substance that decreases reaction rates.
• Inhibitors will be introduced in more detail when studying factors that affect enzyme activity (Chapter 20).
• The characteristic function of many poisons and some medicines is to inhibit one or more enzymes and to decrease the rates of the reactions they catalyze.
• Some substances normally found in cells inhibit specific enzyme-catalyzed reactions and thereby provide a means for the internal regulation of cellular metabolism.

• 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:
H
2
(g) + I
2
(g) 2HI(g)
• When the reaction rate toward the right is equal to the reaction rate toward the left, the reaction is said to be in a state of equilibrium .

• When a reaction is in a state of 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 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.
⇌

• 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:
⇆
• 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.

• The equilibrium expression for the general equation on the previous slide is written as follows:
K

   
    b
X
• 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:
2
⇆ N
2
(g) + 2H
2
O (g)

  
2 2
   
2
2
2

• The K in an equilibrium expression 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.
K

   
    b
X

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 x10 40 , 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.

• According to Le Châtelier's principle , the position of 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.

FACTORS THAT INFLUENCE THE POSITION OF
EQUILIBRIUM (continued)
• If reactants are added, the position of equilibrium will shift toward the products.
• If reactants are removed, the position of equilibrium will shift toward the reactants.
• If products are added, the position of equilibrium will shift toward the reactants.
• If products are removed, the position of equilibrium will shift toward the products.
• When the temperature of a reaction is increased, heat is added.
• If the reaction is endothermic, heat is a reactant and the position of equilibrium will shift toward the products.
• If the reaction is exothermic, heat is a product and the position of equilibrium will shift toward the reactants.

FACTORS THAT INFLUENCE THE POSITION OF
EQUILIBRIUM (continued)
• When the temperature of a reaction is decreased, heat is removed.
• If the reaction is endothermic, heat is a reactant and the position of equilibrium will shift toward the reactants.
• If the reaction is exothermic, heat is a product and the position of equilibrium will shift toward the products.
• Catalysts cannot change the position of equilibrium because they lower 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 .

• Consider the following endothermic reaction at equilibrium: heat + 4NO
2
(g) + 6H
2
O(g) 7O
2
(g) + 4NH
3
(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 NO
2 was added to an equilibrium mixture, the equilibrium position would again shift toward the right to try to use up the added NO
2
.
• If some NH
3 was removed from an equilibrium mixture, the equilibrium position would again shift toward the right in an attempt to replace the NH
3 that was removed.
• If a catalyst were added, the equilibrium position would not change.