Chapter 14: Chemical
Equilibrium
Renee Y. Becker
Valencia Community College
1
Introduction
1. How far does a reaction proceed toward completion
before it reaches a state of chemical equilibrium?
2. Chemical equilibrium
a) The state reached when the concentrations of
reactants and products remain constant over time
b) A state in which the concentration of reactants
and products no longer change (net)
3. Equilibrium mixture
A mixture of reactants and products in the equilibrium
state
2
Introduction
4. What are we interested in?
a) What is the relationship between the
concentration of reactants and products in an
equilibrium mixture?
b) How can we determine equilibrium
concentrations from initial concentrations?
c) What factors can be exploited to alter the
composition of an equilibrium mixture?
3
The Equilibrium State
• In previous chapters we have generally assumed
that chemical reactions result in complete
conversion of reactants to products
Many reactions do not go to completion!!
Example1:
4
The Equilibrium State
5
The Equilibrium State
The two experiments demonstrate that the interconversion
of N2O4 and NO2 is reversible and that the same
equilibrium state is reached starting from either
substance.
1. This is why we use a  instead of 
2. Since both NO2 and N2O5 are products and reactants
we will call the chemical on the left reactants and
on the right products.
3. All chemical reactions are reversible
6
The Equilibrium State
4. We call a reaction irreversible when it proceed
nearly to completion
a. Equilibrium mixture contains almost all
products and almost no reactants
b. Reverse reaction is too slow to be detected
5. In an equilibrium state the reaction does not stop at
particular concentrations of reactants and products,
the rates of the forward and reverse reactions
become equal.
Important: reaction does not stop
7
The Equilibrium State
6. Chemical equilibrium is a dynamic state in which
forward and reverse reactions continue at equal rates so
that there is no net conversion of reactants to products
8
Example 1:
Which of the following is correct?
1. Some reactions are truly not reversible
2. All reactants go to all products in all reactions
3. All reactions are reversible to some extent
4. The rates of the forward and reverse reactions
will never be equal
9
The Equilibrium Constant, Kc
General equation:
aA + bB  cC + dD
Equilibrium equation: Kc = [C]c [D]d
[A]a [B]b
products
reactants
The substances in the equilibrium equation must be
gases or molecules and ions in solution, NO
SOLIDS! NO PURE LIQUIDS!
Kc units are omitted but you must say at what
temperature!
10
The Equilibrium Constant, Kc
a A + bB  cC + dD
Kc = [C]c [D]d
[A]a [B]b
If we write the equation in the reverse direction
cC + dD  aA + bB
K’c = [A]a [B]b = 1
[C]c [D]d
kc
11
Example 2:
Write the equilibrium equation for each of the
following reaction:
a)
b)
N2(g) + 3 H2(g)  2 NH3(g)
2 NH3(g)  N2(g) + 3 H2(g)
12
Example 3:
• The oxidation of sulfur dioxide to give sulfur trioxide is
an important step in the industrial process for synthesis
of sulfuric acid. Write the equilibrium equation for
each of the following reactions:
a)
b)
2 SO2(g) + O2(g)  2 SO3(g)
2 SO3(g)  2 SO2(g) + O2(g)
The following equilibrium concentrations were measured at 800 K:
[SO2] = 3.0 x 10-3 M [O2] = 3.5 x 10-3 M
[SO3] = 5.0 x 10-2 M
Calculate the equilibrium constant at 800 K a and b
13
The Equilibrium Constant Kp
Kp = equilibrium constant with respect to partial
pressures of reactants and products
a A + bB  cC + dD
Kp = (PC)c (PD)d
(PA)a (PB)b
Relationship between Kc and Kp
Kp = Kc(RT)n
14
Example 4:
In the industrial synthesis of hydrogen, mixtures of
CO and H2O are enriched in H2 by allowing the
CO to react with steam. The chemical equation
for this so-called water-gas shift reaction is:
CO(g) + H2O(g)  CO2(g) + H2(g)
What is the value of Kp at 700 K if the partial
pressures in an equilibrium mixture at 700 K are
1.31 atm of CO, 10.0 atm of H2O, 6.12 atm of
CO2, and 20.3 atm H2?
15
Example 5:
When will kc = kp ?
1. 2 SO2(g) + O2(g)  2 SO3(g)
2. CO(g) + H2O(g)  CO2(g) + H2(g)
3. N2(g) + 3 H2(g)  2 NH3(g)
16
Example 6:
Nitric oxide reacts with oxygen to give nitrogen
dioxide, an important reaction in the Ostwald
process for the industrial synthesis of nitric acid:
2 NO(g) + O2(g)  2 NO2(g)
a) If Kc = 6.9 x 105 @ 227C, what is the value of
Kp @ 227C?
b) If Kp = 1.3 x 10-2 @ 1000 K, what is the value of
Kc @ 1000 K?
17
Heterogeneous Equilibria
Introduction
1. So far we have been talking about
homogeneous equilibria, in which all reactants
and products are in a single phase (gas or
solution)
2. Heterogeneous equilibria are those in which
reactants and products are present in more
than one phase
18
Example 7:
For each of the following reactions, write the
equilibrium constant expression for Kc
a) 2 Fe(s) + 3 H2O(g)  Fe2O3(s) + 3 H2(g)
b) 2 H2O(l)  2 H2(g) + O2(g)
c) SiCl4(g) + 2 H2(g)  Si(s) + 4 HCl(g)
d) Hg22+(aq) + 2 Cl-(aq)  Hg2Cl2(s)
19
Example 8:
Which of the following has a Heterogeneous
equilibria?
1. 2 SO2(g) + O2(g)  2 SO3(g)
2. CO(g) + H2O(g)  CO2(g) + H2(g)
3. SiCl4(g) + 2 H2(g)  Si(s) + 4 HCl(g)
20
Using the Equilibrium Constant
Introduction
Knowing the value of the equilibrium constant for a
chemical reaction lets us:
1. Judge the extent of the reaction
2. Predict the direction of the reaction
3. Calculate the equilibrium concentrations from any
initial concentrations
21
Using the Equilibrium Constant
The numerical value of the equilibrium constant for a reaction
indicates the extent to which reactants are converted to
products
1.
Large value for Kc > 103 reaction proceeds essentially to
100% (mostly products)
2.
Small value for Kc < 10-3 reaction proceeds hardly at all before
equilibrium is reached (mostly reactants)
3. If a reaction has an intermediate value of Kc = 103 to 10-3
a. Appreciable concentrations of both reactants and
products are present in the equilibrium mixture
22
Predicting the direction of Reaction
Reaction Quotient = Qc
1. Not necessarily equilibrium concentrations, at some time, t,
snapshot of reaction
2. As time passes, Qc changes toward the value of Kc
3. When the equilibrium state is reached Qc = Kc
4. Qc allows us to predict the direction of reaction by
comparing the values of Kc and Qc
a) If Qc< Kc, net reaction goes from left to
right, (reactant to products)
b) If Qc > Kc, net reaction goes from right to
left, (products to reactants)
c) If Qc = Kc, no net reaction occurs
23
Example 9:
The equilibrium constant for the reaction
2 NO(g) + O2(g)  2 NO2(g)
is 6.9 x 105 @ 500 K. A 5.0 L reaction vessel at this
temperature was filled with 0.060 mol of NO, 1.0 mol
O2, and 0.80 mol NO2.
a) Is the reaction mixture at equilibrium? If not, in which
direction does the net reaction proceed?
b) What is the direction of the net reaction if the initial amounts
are 5.0 x 10-3 mol of NO, 0.20 mol of O2 and 4.0 mol of NO2?
24
Factors that Alter the Composition of an Equilibrium Mixture
Introduction
One of the principal goals of chemical synthesis is to maximize
the conversion of reactants to products while minimizing the
expenditure of energy.
1. Can be achieved if the reaction goes nearly to completion
at mild temperatures and pressures.
2. If the equilibrium mixture is high in reactants and poor in
products, the experimental conditions must be
changed.
3. Several factors can be exploited to alter the composition of
an equilibrium mixture.
A.
The concentration of reactants or products
B.
The pressure and volume
C.
The temperature
25
Le Chatelier’s Principle
Le Chatelier’s Principle
If a stress is applied to a reaction mixture at equilibrium,
net reaction occurs in the direction that relieves the stress
1. Stress means a change in the concentration, pressure,
volume, or temperature that disturbs the original
equilibrium
2. Reaction then occurs to change the composition of
the mixture until a new state of equilibrium is reached
3. The direction that the reaction takes (reactants to
products or products to reactants) is the one that
reduces the stress
26
Altering an Equilibrium Mixture: Changes in Concentration
In general, when an equilibrium is disturbed by the
addition or removal of any reactant or product, Le
Chatelier’s principle predicts that:
1. The concentration stress of an added reactant or
product is relieved by net reaction in the direction
that consumes the added substance
2. The concentration stress of a removed reactant or
product is relieved by net reaction in the direction
that replenishes the removed substance
27
Example 10:
Consider the equilibrium for the water-gas shift reaction:
CO(g) + H2O(g)  CO2(g) + H2(g)
Use Le Chatelier’s principle to predict how the
concentration of H2 will change and what direction the
reaction will flow when the equilibrium is disturbed by:
1.
2.
3.
4.
Adding CO
Adding CO2
Removing H2O
Removing CO2
28
Example 11:
In the following reaction, if I take away CO, which
direction will the reaction proceed to equilibrium?
CO2(g) + H2(g) CO(g) + H2O(g)
1. Products 
2. Reactants 
29
Altering an Equilibrium Mixture: Changes in Pressure and Volume
In general Le Chatelier’s Principle predicts that:
1. An increase in pressure by reducing the
volume will bring about net reaction in
the direction that decreases the number
of moles of gas
2. A decrease in pressure by enlarging the
volume will bring about net reaction in
the direction that increases the number
of moles of gas.
30
Example 12:
Which direction will the reaction flow if the
following equilibria is subjected to an increase in
pressure by decreasing the volume?
1.
CO(g) + H2O(g)  CO2(g) + H2(g)
2.
2 CO(g)  C(s) + CO2(g)
3.
N2O4(g)  2 NO2(g)
31
Example 13:
If I increase the pressure by decreasing the volume,
which direction will the reaction flow to reach
equilibrium?
C(s) + CO2(g)  2 CO(g)
1. Products 
2. Reactants 
32
Altering the Equilibrium Mixture: Changes in Temperature
In general, the temperature dependence of the equilibrium constant
depends on the sign of H for the reaction
1.
The equilibrium constant for an exothermic reaction
(negative H) decreases as the temperature increases
2.
The equilibrium constant for an endothermic reaction
(positive H) increases as the temperature increases.
3.
H = standard enthalpy of reaction, enthalpy change
measured under standard conditions
4.
Standard conditions = most stable form of a substance at 1
atm pressure and at a specified temperature, usually 25C; 1
M concentration for all substances
33
Altering the Equilibrium Mixture: Changes in Temperature
Le Chatelier’s Principle says that if heat is added to an equilibrium
mixture (increasing the temperature) net reaction occurs in the
direction that relieves the stress of the added heat.
1.
For an endothermic reaction heat is absorbed by reaction in
the forward direction. The equilibrium shifts to the right at
the higher temperatures, Kc increases with increasing
temperature
2.
For an exothermic heat is absorbed by net reaction in the
reverse direction, so Kc decreases with temperature, and the
reaction would flow to the left (reactants)
34
Example 14:
When air is heated at very high temperatures in an
automobile engine, the air pollutant nitric oxide is
produced by the reaction
N2(g) + O2(g)  2 NO(g)
H = 180.5 kJ
1. How does the equilibrium amount of NO
vary with an increase in temperature?
2. What direction is the net reaction flowing?
35
The Effect of a Catalyst on Equilibrium
A catalyst increases the rate of a chemical reaction by making
available a new, lower-energy pathway for conversion of
reactants to products.
1.
Since the forward and reverse reaction pass through the same
transition state, a catalyst lowers the activation energy for
both
2.
The rates of the forward and reverse reactions increase by the
same factor
3.
Catalyst accelerates the rate at which equilibrium is reached
4.
Catalyst does not affect the composition of the equilibrium
mixture
36
The Effect of a Catalyst on Equilibrium
37
Example 15:
A platinum catalyst is used in automobile catalytic converters to
hasten the oxidation of carbon monoxide:
2 CO(g) + O2(g)  2 CO2(g)
H = -566 kJ
Suppose that you have a reaction vessel containing an
equilibrium mixture. Will the amount of CO increase,
decrease, or remain the same when:
1.
2.
3.
4.
5.
A platinum catalyst is added
The temperature is increased
The pressure is increased by decreasing the volume
The pressure is increased by adding argon gas
The pressure is increased by adding O2 gas
38
The Link Between Chemical Equilibrium and Chemical Kinetics
A + B  C + D
Assuming that the forward and reverse reactions occur in a single
bimolecular step, elementary reactions, we can write the
following rate laws
Rate of forward reaction = kf [A] [B]
Rate of reverse reaction = kr [C] [D]
When t=0 [C] = [D] = 0
As A and B are converted to C and D the rate of the forward
reaction decreases and the rate of the reverse reaction is
increasing, until they are equal, chemical equilibrium
kf [A] [B] = kr [C] [D]
39
The Link Between Chemical Equilibrium and Chemical Kinetics
kf = [C] [D]
kr
[A] [B]
The right side of this equation is the equilibrium constant
expression for the forward reaction, which equals the
equilibrium constant Kc
Kc = [C] [D]
[A] [B]
Therefore the equilibrium constant is simply the ratio of
the rate constants for the forward and reverse reactions:
Kc = kf
kr
40
Example 16:
Nitric oxide emitted from the engines of supersonic transport planes can
contribute to the destruction of stratospheric ozone:
NO(g) + O3(g)  NO2(g) + O2(g)
This reaction is highly exothermic (E = -200 kJ), and its equilibrium
constant Kc is 3.4 x 1034 at 300 K
1.
Which rate constant is larger, kf or kr?
2.
The value of kf at 300 K is 8.5 x 106 M-1 s-1. What is the value of kr at
the same temperature?
3.
A typical temperature in the stratosphere is 230 K. Do the values of kf,
kr, and Kc increase or decrease when the temperature is lowered from
300 K to 230 K?
41
Example 17:
If I increase the temperature of reaction which
way will the reaction flow to equilibrium?
NO2(g) + O2(g)  NO(g) + O3(g)
H = 200 kJ
1. Products 
2. Reactants 
42