Chapter 15. Chemical Equilibrium
15.1 The Concept of Equilibrium
Consider colorless frozen N2O4.
At room temperature, it decomposes to brown NO2.
N2O4(g) → 2NO2(g)
At some time, the color stops changing and we have a mixture of N2O4 and NO2.
Dynamic Chemical equilibrium is the point at which the concentrations of all species are constant.
Assume that both the forward and reverse reactions are elementary processes.
We can write rate expressions for each reaction.
Forward reaction:
N2O4(g) → 2ΝΟ2(g)
Ratef = kf[N2O4]
kf = rate constant (forward reaction)
Reverse reaction:
2NO2(g) → N2O4(g)
2
Rater = kr[NO2]
kr = rate constant (reverse reaction)
Place some pure N2O4 into a closed container.
As N2O4 reacts to form NO2, the concentration of N2O4 will decrease and the concentration of NO2 will
increase. Thus, we expect the forward reaction rate to slow and the reverse reaction rate to increase.
Eventually we get to equilibrium where the forward and reverse rates are equal.
At equilibrium:
kf[N2O4] = kr[NO2]2
Rearranging, we get:
kf/kr = a constant = Keq
At equilibrium the concentrations of N2O4 and NO2 do not change.
This mixture is called an equilibrium mixture.
This is an example of a dynamic equilibrium.
A dynamic equilibrium exists when the rates of the forward and reverse reactions are equal.
No further net change in reactant or product concentration occurs.
The double arrow ↔ implies that the process is dynamic.
Chemical Equilibrium page 2
15.2 The Equilibrium Constant
Consider the Haber process:
N2(g) + 3H2(g) ↔ 2NH3(g)
It is used for the preparation of ammonia from nitrogen and hydrogen.
The process is carried out at high temperature (500˚C) and pressure (200 atm) in the presence of
a catalyst.
Ammonia is a good source of fixed nitrogen for plants.
Much of the NH3 produced industrially is used as a fertilizer.
If we start with a mixture of nitrogen and hydrogen (in any proportions), the reaction will reach
equilibrium with constant concentrations of nitrogen, hydrogen, and ammonia.
However, if we start with just ammonia and no nitrogen or hydrogen, the reaction will proceed and N2
and H2 will be produced until equilibrium is achieved.
No matter what the starting compositions of reactants and products are, the equilibrium mixture contains
the same relative concentrations of reactants and products.
Equilibrium can be reached from either direction!
We can write an expression for the relationship between the concentration of the reactants and products
at equilibrium.
This expression is based on the law of mass action.
For a general reaction,
aA + bB ↔ dD + eE
The equilibrium expression is given by:
Kc =
[ D d ][ E ] e
[ A] a [ B ]b
 Where Kc is the equilibrium constant, Keq.
 The subscript “c” indicates that molar concentrations were used to evaluate the constant.
 Note that the equilibrium constant expression has products in the numerator and reactants in
the denominator.
The value of Kc does not depend on initial concentrations of products or reactants.
Consider the reaction:
N2O4(g) ↔ 2NO2(g)
The equilibrium constant is given by:
Kc =
[ NO2 ] 2
[ N 2 O4 ]
The value of this constant (at 100˚C) is 6.50 (regardless of the initial concentrations of N2O4(g)
or NO2(g).
The equilibrium expression depends on stoichiometry.
 It does not depend on the reaction mechanism.
 The value of Kc varies with temperature. Why?
We generally omit the units of the equilibrium constant. However, it is useful to use it when we our
trying to evaluate concentration at equilibrium; a concentration term on the right hand side of the
equilibrium expression.
Equilibrium Constants in Terms of Pressures, Kp.
When the reactants and products are gases we can write an equilibrium expression using partial
pressures rather than molar concentrations.
The equilibrium constant is Kp where “p” stands for pressure.
Chemical Equilibrium page 3
For the reaction:
aA + bB  dD + eE
Kp =
( PD ) d ( PE ) e
( PA ) a ( PB ) b
The numerical values of Kc and Kp will differ.
They can be interconverted using the ideal gas equation and our definition of molarity:
PV = nRT so P = (n/V)RT
If we express volume in liters the quantity (n/V) is equivalent to molarity.
Thus, the partial pressure of a substance, A, is given as:
P = [A]RT
We can use this to obtain a general expression relating Kc and Kp:
Kp = Kc(RT)∆n
where ∆n = (moles of gaseous products) – (moles of gaseous reactants).
15.3 Interpreting and Working with Equilibrium Constants
The Magnitude of Equilibrium Constants
The equilibrium constant, K, is the ratio of products to reactants.
Therefore, the larger K is, the more products are present at equilibrium.
Conversely, the smaller K is, the more reactants are present at equilibrium.
If K >> 1, then products dominate at equilibrium and equilibrium lies to the right.
If K << 1, then reactants dominate at equilibrium and the equilibrium lies to the left.
The Direction of the Chemical Equation and its corresponding Keq.
As mentioned earlier, the equilibrium can be approached from either direction.
Consider the reaction:
N2O4(g) ↔ 2NO2(g)
The equilibrium constant for this reaction (at 100˚C) is:
Kc =
[ NO2 ] 2
= 0.212
[ N 2 O4 ]
However, when we write the equilibrium expression for the
reverse reaction,
2NO2(g) ↔ N2O4(g)
The equilibrium constant for this reaction (at 100˚C) is:
[N2O4 ]
= 4.72
[NO2 ]2
The equilibrium constant for a reaction in one direction is
K 'c =
€
Chemical Equilibrium page 4
the reciprocal of the equilibrium constant of the reaction in the opposite direction.
1
K 'c
Relating Chemical Equations and Equilibrium Constants
 The equilibrium constant of a reaction in the reverse direction is the inverse of the equilibrium
constant of the reaction in the forward
€ direction.
 The equilibrium constant of a reaction that has been multiplied by a number is the equilibrium
constant raised to a power equal to that number.
 The equilibrium constant for a net reaction made up of two or more steps is the product of the
equilibrium constants for the individual steps.
Kc =
15.4 Heterogeneous Equilibria
Equilibria in which all reactants and products are present in the same phase are called homogeneous
equilibria. Equilibria in which one or more reactants or products are present in a different phase are
called heterogeneous equilibria.
Consider the decomposition of calcium carbonate:
CaCO3(s) ↔ CaO(s) + CO2(g)
Experimentally, the amount of CO2 does not depend on the amounts of CaO and CaCO3.

Why?
The concentration of a pure solid or pure liquid equals its density divided by its molar mass.
[CaO][CO 2 ]
[CaCO 3 ]
Neither density nor molar mass is a variable!
Thus, the concentrations of solids and pure liquids are constant.
For the decomposition of CaCO3:
CaO and CaCO3 are€pure solids and have constant concentrations.
Kc =
Kc =
(constant 1)[CO 2 ]
(constant 2)
Rearranging:
∴Kc '= Kc
(constant 2)
= [CO2 ]
(constant 1)
K p = PCO2
 If a pure solid or a pure liquid is involved in a heterogeneous equilibrium, its concentration is not
included in the equilibrium constant expression.
 Therefore, we anticipate that the amount of CO2 formed will not depend on the amounts of CaO
and CaCO3 present.
Note that although the concentrations of these species are not included in the equilibrium
expression, they do participate in the reaction and must be present for an equilibrium to be
established!
Other common examples of heterogeneous equilibria include:
 Dissolution of a solid to form a saturated solution.
Example:
PbCl2(s) ↔ Pb2+(aq) + 2Cl–(aq)
Ksp= Kc[PbCl2] = [Pb2+][Cl–]2
Here “sp” in Ksp, stands for solubility products constant.
Chemical Equilibrium page 5
 Systems where the solvent is involved as a reactant or product and the solutes are present at low
concentrations.
Example:
H2O(l) + CO32–(aq) ↔ OH–(aq) + HCO3–(aq)
−
[OH− ][HCO3 ]
[CO3 2− ]
Here the concentration of water, the solvent, is essentially constant and we can think of it as a pure
liquid. The concentration of water, [H2O], is incorporated into the equilibrium constant as shown above.
K'c = Kc [H2O] =
15.5 Calculating€Equilibrium Constants
Proceed as follows:
Tabulate initial and equilibrium concentrations (or partial pressures) for all species in the
equilibrium reaction.
If an initial and an equilibrium concentration is given for a species, calculate the change in
concentration. Use the coefficients in the balanced chemical equation to calculate the changes in
concentration of all species.
Deduce the equilibrium concentrations of all species.
Use the equilibrium concentrations to calculate the value of the equilibrium constant.
(Examples will follow to clarify what is summarized here and below in section 15.6)
15.6 Applications of Equilibrium Constants
Predicting the Direction of Reaction
 For a general reaction:
aA + bB ↔ dD + eE
 We define Q, the reaction quotient, as:
Q=
[D]d [E]e
[A]a [B]b
 Where [A], [B], [D], and [E] are molarities (for substances in solution) or partial pressures
(for gases) at any given time.
 We can compare Qc with Kc or Qp with Kp:
 If Q = K, then the system is at equilibrium, no further net change occurs.
 If Q < K, then the forward reaction must occur; decreasing the concentration of the
reactants and increasing the concentration of the products until equilibrium is
reached. Q increases until it equals K.
 If Q > K, then the reverse reaction must occur to reach equilibrium. Products are
consumed and reactants are formed. Q decreases until it equals K.
Calculating Equilibrium Concentrations
The same steps used to calculate equilibrium constants are used to calculate equilibrium concentrations.
Generally, we do not have a number for the change that occurs in concentration.
 Therefore, we need to assume that x mol/L of a species is produced (or used).
The equilibrium concentrations are given as algebraic expressions, to be evaluated.
Calculating Equilibrium Constant.
Example 1:
Chemical Equilibrium page 6
Consider the following equilibrium reaction:
E
+ F
↔ G + H
in a liter vessel.
initially(i):
5.0 mol 7.0 mol
--at equilibrium(e):
4.5 mol
of G is found.
What is Kc =?
In order to answer this question we must consider the following:
(a)
what is initially present
(b)
How much is consumed by the reaction, or
(c)
How much is produced by the reaction, and
(d)
what is present at equilibrium.
First we set up our ice table; initial concentrations, change, and equilibrium concentrations.
{We will complete this example in Lecture!}
E
+ F
↔ G + H
in a liter vessel.
i:
5.0 mol 7.0 mol
--c:
<= from the stoichiometry!
e:
Example 2:
i:
2 HY
1.0 M
↔
H2 + 2 Y
Kc= 4.0 x 10-9 M
Calculate [Y]equil
Example 3: 20.0 g of CaCO3(s) is allowed to reach equilibrium in a 10.0 liter vessel at 800 ˚C, at
which Kp = 1.16 atm. Determine the %CaCO3 remaining at equilibrium. The equilibrium equation is
20.0 g
CaCO3(s)
↔ CaO(s) + CO2(g)
Example 4: Nitrosyl bromide, NOBr(g) is 34% dissociated into NO(g) + Br2(g) at 25.0˚C. The total
pressure is equal to 0.25 atm. Calculate Kp and Kc.
Chemical Equilibrium page 7
i:
2 NOBr(g)
x atm
↔
2 NO(g)
+
Br2(g)
15.7 Le Châtelier’s Principle
Consider the Haber process again:
N2(g) + 3H2(g) ↔ 2NH3(g)
 It is known that as the pressure increases, the amount of ammonia present at equilibrium
increases.
 As the temperature increases, the amount of ammonia at equilibrium decreases.
 Can this be predicted?
Yes it Can! The solution to this question is expressed in Le Châtelier’s principle: If a stress is
applied to a system at equilibrium (by a change in temperature, a change in pressure, or a change in the
concentration of one or more components), the system will shift its equilibrium position in a direction to
relieve the stress (counteract the effects of the stress) and return to equilibrium unaided.
Change in Reactant or Product Concentration
If a chemical system is at equilibrium and we add or remove a product or reactant, the reaction will shift
so as to reestablish equilibrium.
For example, consider the Haber process again:
N2(g) + 3H2(g) ↔ 2NH3(g)
 If H2 is added to the system at equilibrium, the reaction quotient, Q < K.
 The system must respond to counteract the added H2 (Le Châtelier’s principle).
To relieve the stress of the added H2, the system must consume the H2. It does this by
reacting with N2 to produce more products until equilibrium is reestablished.
Therefore, [H2] and [N2] will decrease and [NH3] will increase until Q = K.
We can exploit this industrially.
Suppose that we wanted to optimize the amount of ammonia we formed from the Haber process.
 We might flood the reaction vessel with reactants and continuously remove product.
 The amount of ammonia produced is optimized because the product (NH3) is
continuously removed and the reactants (N2 and H2) are continuously being added.
Effects of Volume and Pressure Changes
Consider a system at equilibrium.
If the equilibrium involves gaseous products or reactants, the concentration of these species will be
changed if we change the volume of the container.
 For example, if we decrease the volume of the container, the partial pressures of each gaseous
species will increase. Why?
Chemical Equilibrium page 8
 Le Châtelier’s principle predicts that if pressure is increased, the system will shift to counteract
the increase.
That is, the system shifts to decrease the number of moles of gases and decrease pressure.
Therefore, an increase in pressure favors the direction that has fewer moles of gas.
Consider the following system:
N2O4(g) ↔ 2NO2(g)
An increase in pressure (by decreasing the volume) favors the formation of colorless N2O4.
The instant the pressure increases, the concentration of both gases increases and the system is not
at equilibrium. The stress placed on the system is an increase in pressure!
The system shifts to reduce the increase pressure by reducing the number moles of gas.
The partial pressures adjust until equilibrium is reestablished.
The mixture will be lighter in color. Some of the brown NO2 has been converted into colorless
N2O4(g)
 In a reaction with the same number of moles of gas in the products and reactants, changing the
pressure has no effect on the equilibrium.
 Similarily, no change will occur if we increase the total gas pressure by the addition of a gas
that is not involved in the equilibrium reaction.
Effect of Temperature Changes
The equilibrium constant is temperature dependent! Why?
A change in temperature will give a new equilibrium constant. The question is how will a change in
temperature alter a system at equilibrium?
 It depends on whether the particular reaction is endothermic or exothermic.
For example, consider the endothermic reaction:
Co(H2O)62+(aq) + 4Cl–(aq) ↔ CoCl42–(aq) + 6H2O(l)
∆H > 0
Co(H2O)62+ is pale pink and CoCl42– is a deep blue.
Consider the following experiments:
At room temperature, an equilibrium mixture (light purple) is placed in a beaker of warm water.
 The mixture turns deep blue.
 This indicates a shift toward products (blue CoCl42–).
 This reaction is endothermic.
 For an endothermic reaction (∆H > 0), heat can be considered as a reactant!
Thus, adding heat causes a shift in the forward direction.
The room-temperature equilibrium mixture is placed in a beaker of ice water.
 The mixture turns bright pink.
 This indicates a shift toward reactants (pink Co(H2O)62+).
 In this case, by cooling the system we are removing a reactant (heat).
 Thus, the reaction is shifted in the reverse reaction.
A change in temperature causes a change in the value of K! (Why?)
If we increase the temperature of an endothermic reaction, K increases.
If we increase the temperature of an exothermic reaction, K decreases.
The Effect of Catalysts
A catalyst lowers the activation energy barrier for the reaction.
Therefore, a catalyst will decrease the time taken to reach equilibrium.
A catalyst does not effect the composition of the equilibrium mixture.
Example #5: Consider the following system in equilibrium
Chemical Equilibrium page 9
N2(g) + 2 O2(g) ↔ 2 NO2(g)
∆H˚= 66.4 kJ/mol
What effect (shift right, shift left, no change) on the system in equilibrium will the following change
have?
Change
Effect
(a)
add more N2
(b)
increase volume of container
(c)
remove O2
(d)
add catalyst
(e)
add Ne gas
(f)
decrease T
Example #6: The four substances HCl, I2, HI, and Cl2 are mixed in a reaction vessel and allowed to
reach equilibrium. Certain changes (which are specified in the first column in the table below) are
then made to this mixture. Considering each change separately, and with the aid of Le Chatelier's
principle predict the effect ([CIRCLE], increase, I; decrease, D; or no change, NC) that the change
has on the original equilibrium value of the quantity in the second column of the table for the
following equilibrium system:
2 HCl(g) + I2(s) ↔ 2 HI(g) + Cl2(g)
∆H˚= +237.6 kJ
Change
decrease volume
increase volume
raise temperature
add I2
add HCl
remove HI
increase pressure
lower pressure
add I2
add Cl2
Quantity
Kc
amount of Cl2
Kc
amount of Cl2
amount of HI
amount of Cl2
amount of HCl
amount of Cl2
Kc
concentration of I2
Effect
I
I
I
I
I
I
I
I
D
D
D
D
D
D
D
D
NC
NC
NC
NC
NC
NC
NC
NC
I
I
D
D
NC
NC