18.2 Reversible Reactions and Equilibrium
Standards: 8b, 9a, 9b, 9c
Vocabulary
reversible reaction, chemical equilibrium, equilibrium position, Le Chatelier’s principle,
equilibrium constant
Connection to your world
For years, scientist tried to produce nitrogen compounds for fertilizers to increase the amount of available
foodstuffs. Unfortunately, none of these efforts proved to be commercially successful. Finally, in the early
1900s, two German chemists, Fritz Haber and Karl Bosch, refined the process of making ammonia from
elemental nitrogen and hydrogen. Their success came from controlling the temperature and pressure under
which two gases are reacted. In this section, you will learn how changing the reaction conditions can influence
the yield of a chemical reaction.
Reversible Reaction
From what you have already learned, you may have inferred that chemical reaction go completely from
reactants to products, just as the chemical equation indicates this is not usually the case. Some reactions are
reversible. A reversible reaction is one in which the conversion of reactants to products and the conversion of
products to reactants occur simultaneously. One example of a reversible reaction is the following.
Forward reaction: 2SO2(g) + O2(g)  2SO3(g)
Reverse reaction: 2SO2(g) + O2(g)  2SO3(g)
In the first section, which is read from left to right, sulfur dioxide and oxygen produce sulfur trioxide? In the
second, which id read from right to left, sulfur trioxide decomposes into oxygen and sulfur dioxide. The first
section is called the forward reaction. The second is called the reverse reaction. The two equations can be
combined into one using a double arrow. The double arrow tells you that this reaction is reversible.
2SO2(g) + O2(g)
Sulfur
dioxide
Oxygen
PCl5(g) + heat
2SO3(g)
Sulfur
trioxide
PCl3(g) + Cl2(g)
These graphs show how the concentrations of A, B, C, and D over time
What actually happens when sulfur dioxide and oxygen gases are mixed in a sealed chamber? The two reactants
begin to react to form sulfur trioxide at a particular rate. Because no sulfur trioxide is present at the beginning of
the reaction, the initial rate of the reverse reaction is zero. As sulfur trioxide begins to form, however, the
decomposition of sulfur trioxide begins. This reverse reaction proceeds slowly at first, but its rate increased as
the concentration of sulfur trioxide increase. Simultaneously, the rate of the forward reaction decreases because
sulfur dioxide and oxygen are begin used up. Eventually sulfur trioxide is decomposing to sulfur trioxide and
oxygen as fast as sulfur dioxide and oxygen are forming sulfur trioxide. When the rates of the forward and
reverse reaction are equal, the reaction has reached a state of balance called chemical equilibrium. Changes in
concentrations of the three components during the course of the reaction are shown in the graphs in figure
18.10. The graph on the left shows the process of the reaction that start with specific concentration of SO2 and
O2 but with zero concentration of SO3. The graph on the right shows concentration for a reaction that begins
with an initial concentration of SO3 and zero concentration for SO2 and O2 . Notice that after a certain time, all
concentration remain constant. At chemical equilibrium, no net change occur s in the actual amounts of the
components of the system. The amount of SO3 in the equilibrium mixture is the maximum amount that can be
produced by this reaction under the conditions of the reaction.
The unchanging amounts of SO2, O2 and SO3 in the reaction mixture at equilibrium might cause you to think
both reactions have stopped. This is not the case. Chemical equilibrium is a dynamic state. Both the forward and
reverse reactions continue, but because their rates are equal, no net change occur in their concentrations. The
escalators in Figure 18.11 are like the double arrows in a dynamic equilibrium equation. The number of the
people using the up escalator must equal the number of the people using the down escalator for the number of
the people on both floors to remain constant
√Checkpoint: Why is chemical equilibrium called a dynamic state?
Although the rates of the forward and reverse reactions are equal at chemical equilibrium, the
concentrations of the components on both sides of the chemical equation are not necessarily the same. In fact,
they can be dramatically different. Figure 18.10 shows the equilibrium concentrations of SO2, O2, and SO3.
The concentration of SO3 is significantly greater then the concentration of SO2, and O2. The relative
concentrations of the reactants and products at equilibrium constitute the equilibrium position of a reaction.
The equilibrium position indicates whether the reactants or products are favored in a reversible reaction. If A
reacts to give B and the equilibrium mixture contains significantly more of B—say 1% A and 99% B—then the
formation of B is said to be favored.
A
B
1%
99%
On the other hand, if the mixture contains 99% A and 1% B at equilibrium, then the formation of A is favored.
A
99%
B
1%
Notice that the equilibrium arrows are not of equal length; the longer of the two arrows indicates the favored
direction of the reaction.
In principle, almost all reactions are reversible to some extent under the right conditions. In practice, one
set of components is often so favored at equilibrium that the other set cannot be detected. If one set of
components (reactants) is completely converted to new substances (products), you can say that the reaction has
gone to completion, or is irreversible. When you mix chemicals expecting to get a reaction but no products can
be detected, you can say that there is no reaction. Reversible reactions occupy a middle ground between the
theoretical extremes of irreversibility and no reaction.
A catalyst speeds up both the forward and the reverse reactions equally because the reverse reaction is
exactly the opposite of the forward reaction. The catalyst lowers the activation energy of the reaction by the
same amount in both the forward and reverse directions. Catalysts do not affect the amounts of reactants and
products present at equilibrium; they simply decrease the time it takes to establish equilibrium.
√Checkpoint: What is chemical equilibrium?
Factors Affecting Equilibrium: Le Chatelier’s principle
A delicate balance exists in a system at equilibrium. Changes of almost any kind can disrupt this balance. When
the equilibrium of a system is disrupted, the system makes adjustments to restore equilibrium. However, the
equilibrium position of the restored equilibrium is different from the original equilibrium position; that is, the
amount of products ad reactants may have increased or decreased. Such a change is called a shift in the
equilibrium position.
The French chemist Henri Le Chaterlier (1850-1936) study how the equilibrium position shifts as a
result of changing conditions. He proposed what has come to be called Le Chatelier’s Principle: if a stress is
applied to a system in dynamic equilibrium, the system changes in a way tat relives the stress. Stresses that
upset the equilibrium of a chemical system include changes in the concentration of reactants or products,
changes in temperature, and changes in pressure.
The following examples of applications of Le Chatelier’s principle all involve reversible reactions. For
simplicity and clarity, the components to the left of the reaction arrow will be considered the reactants and the
components to the right of the reaction arrow will be considered the products. Blue arrows indicate the shifts
resulting for additions to or removals from the system. The arrows always point in the direction of the resulting
shift in the equilibrium position-that is, toward the favored side.
Concentration Changing the amount, or concentration, of any reactant or product in a system at equilibrium
disturbs the equilibrium. The system adjusts to minimize the effects of the change. Consider the decomposition
of carbonic acid (h2co3) in aqueous solution to for the products carbon dioxide and water. The system has
reached equilibrium. At equilibrium the amount of carbonic acid is less than 1%.
Add CO2
 Direction of shift
H2CO3(aq)
CO2(aq) + H2O (l)
Remove CO2
Direction shift 
<1%
>99%
Adding more carbon dioxide disrupts the equilibrium. It increases the concentration of CO2 in the mixture and
cause the rate of the reverse reaction to increase. As more reactant (H2CO3) is formed, the rate of the forward
reaction also begins to increase. In time, the rates of the forward and reverse reaction again become equal, and a
new equilibrium is established with a higher concentration of reactant (H2CO3). Adding a product to a reaction
at equilibrium pushes a reversible reaction in the direction of reactants.
If, on the other hand, carbon dioxide is remove, the concentration of CO2 decreases. This causes the rate
of the reverses reaction to decrease. As less reactant (H2CO3) is being formed, the rated of the forward reaction
also begins to decrease. When the rates of the forward and reverse reaction again become equal, equilibrium is
restored but as different equilibrium position. Removing a product always pushes a reversible reaction in the
direction of products.
Farmers use this technique to increase the yield of eggs laid by hens. Hens lay eggs and then proceed to
hatch them. If the eggs are removed after they are laid (removing the product), the hen will lay more eggs
(increasing the yield). Similarly as products are removed from a reaction mixture, they system continually
changes to restore equilibrium by producing more products. But because the products are being removed, the
reactants can never reestablish equilibrium. The reaction continues to produce products until the reactants are
used up. Another example of this concept is found in the body. Blood contains dissolved carbonic acid in
equilibrium with carbon dioxide and water. The body uses the removal of products to keep the concentration of
carbonic acid within a safe range. When the athletes exhale carbon dioxide, the equilibrium shifts toward
carbon dioxide and water, thus reducing the amount of carbonic acid. The same principle applies to adding or
removing reactants. When a reactant is added to a system at equilibrium, the reaction shirts in the direction of
the formation of products. When a reactant is removed, the reaction shifts in the direction of formation of
reactants.
√Checkpoint According to Le Chatelier’s principle, how does a system at equilibrium respond to a
stress?
Temperature Increasing the temperature causes the equilibrium position of a reaction to shift in the direction
that absorbs heat. The heat absorption reduces the applied temperature stress. For example, consider the
following exothermic reaction that occurs when SO3 and O2.
Add heat
 Direction of shift
2SO2 (g) + O2
2SO3 (g) + heat
Remove heat (cool)
Direction of shift 
Heat can be considered to be a product, just like SO3. Heating the reaction mixture at equilibrium pushes the
equilibrium position to the left, which favors the reactants. As a result, the product yield decreases. Cooling, or
removing heat, pulls the equilibrium to the right, and the product yield increases.
Pressure A change in the pressure on a system affects only gaseous equilibria that have unequal number of
moles of reactants and products. An example is the reaction you read about is hydrogen and nitrogen react too
form ammonia. Imagine that the three gases are at equilibrium in a cylinder that has a piston attached to a
plunger-similar to a bicycle pump by with the hose sealed. A catalyst has been including to speed up the
reaction. What happens to the pressure when you push the plunger down? The pressure on the gas
momentarily increases because the same number of molecules is contained in a smaller volume. The system
immediately relieves some of the pressure increase by reducing the number of gas molecules. For every two
molecules of ammonia made, four molecules of reactant are used up (three molecules of hydrogen and one of
nitrogen). Therefore, the equilibrium position shifts to make more ammonia. There are then fewer molecules in
the system. The pressure decreases, although it will not decrease all the way to the original pressure. As you
can see, increasing the pressure on the system results in a shift in the equilibrium position that favors the
formation of product.
Increase pressure
Direction of shift 
N2(g) + 3H2
2NH3 (g)
Reduce pressure
 Direction of shift
The equilibrium position for this reaction can be made to factor the reactants instead of the product. Imagine
pulling the plunger of the piston device back up so the volume containing the gases increases. The increase in
volume decreases the pressure on the system. To restore the higher starting pressure, the system can produce
more gas molecules by decomposition of some ammonia molecules. Decomposition of two molecules of
gaseous NH3 produces four molecules of reactants (three H2 and one N2). Pressure at the new equilibrium is
higher than when the pressure was first decreased, but not as high as it was at the starting equilibrium.
Lowering the pressure on the system thus results in a shift of the equilibrium to favor the reactants.
Conceptual Problem 18.1
Appling Le Chatelier's Principle
What effect do each of the folloing changes have on the equilibrium
position for this reversible rations?
PCl5(g) + heat
a. addition of Cl2
c. Removal of heat
PCl3(g) + Cl2(g)
b. Increase in pressure
c. Removal of PCl3(g) as it is formed
1. Analyze Identify the relevant concepts.
a.-d. The stress placed on each system is
known. The effect of each stress is unknown.
By Le Chatelier's principle, the equilibrium
system will shift in a direction that minimizes
the imposed stress. Analyze the effect of each
change on the reaction
2. Solve Apply concepts to this situation
a. The addition of Cl2, a product, shifts the
b. The equation shows 2 mol of gaseous product
and 1 mol of gaseous reactant. The increase
in pressure is relieved if the equilibrium shifts
to the left, because a decrease in the number
of moles of gaseous substances produces a
decrease in pressure
c. The removal of heat causes the equilibrium
to shift ot the left, because the reverse reaction is heat-producing.
d. The removal of PCl3 causes the equilibrium
equilibrium to the left, forming more PCl5.
to shift to the right to produce more PCl3.
Practice Problems
6. How is the equilibrium position of this reaction
affected by the following changes?
C(s) + H2O(g)
a.
b.
c.
d.
CO(g) + H2(g)
Lowering the temperature
Increasing the pressure
Removing hydrogen
Adding water vapor
CONCEPTUAL PROBLEM 18.2
1.
2.
3.
4.
5.
How do the amounts of reactants and products change after a reaction has reached chemical
equilibrium?
What are three stresses that can upset the equilibrium of a chemical system?
What does the value of the equilibrium constant tell you about the amounts of reactants and products
present at equilibrium?
How can a balanced chemical equation be used to write an equilibrium-constant expression?
Can a pressure change shift the equilibrium position in every reversible reaction? Explain.