CHEMICAL EQUILIBRIA
Many chemical reactions involve an equilibrium process. During an dynamic
equilibrium, the rate of the forward reaction equals the rate of the reverse reaction. This means
that both reactants and products will be present at any given point in time. The equilibrium may
favor either the reactants or products. The extent to which the reaction proceeds towards
products is measured by an equilibrium constant. This constant is a specific ratio of the products
to the reactants. This ratio is often referred to as a mass action expression. Let's look at a
"generic" equilibrium reaction and its mass action expression.
aA + bB  cC + dD
where
A and B are reactants
C and D are products
a,b,c,d are the coefficients in the balanced chemical equation
The mass action expression consists of the product of the products, each raised to the
power given by the coefficient in the balanced chemical equation, over the product of the
reactants, each raised to the power given by the coefficient in the balanced chemical equation.
This mass action expression is set equal to the equilibrium constant, Keq:
Keq 
[C]c [D] d
[A]a [B] b
There are numerous types of equilibrium problems one may encounter, both
qualitative and quantitative. In the subsequent pages we will look at these various types of
problems.
Qualitative Problems
Qualitatively, one may look at how far an equilibrium lies towards the right (towards
products) or left (towards reactants). The magnitude of the equilibrium constant gives one a
general idea of whether the equilibrium favors products or reactants. If the reactants are favored,
then the denominator term of the mass action expression will be larger than the numerator term,
and the equilibrium constant will be less than one. If the products are favored over the reactants,
then the numerator term will be larger than the denominator term, and the equilibrium constant
will be greater than one. The following two equilibria, demonstrate a case in which the reactants
are favored and a case in which the products are favored. Note the magnitude of the equilibrium
constant for each example:
COCl2(g)  CO(g) + Cl2(g)
Keq = 2.2 x 10-10 @ 298 K
(reactant favored)
2NO2(g)  N2O4(g)
Keq = 2.15 x 102 @298 K
(product favored)
An equilibrium is able to shift in either direction, either towards reactants or towards
products, when a stress is placed on the equilibrium. This stress may involve increasing or
decreasing the amount of a reactant or product, changing the volume or pressure in gas phase
equilibria or changing the temperature of the equilibrium system. The direction in which the
equilibrium shifts may be predicted by using Le Chatlier's Principle. Le Chatlier's Principle
states that if a stress is placed on an equilibrium, the equilibrium will shift in the direction which
relieves the stress. Let's look at an equilibrium and the direction it will shift due to various
stresses placed on it.
2 N2 (g) + 3 H2 (g)  2 NH3 (g)
Suppose additional nitrogen (N2) was added to the system. First, identify the stress. The
stress is too much nitrogen. The equilibrium will shift to the right, towards products, in order to
remove the excess nitrogen.
Suppose some ammonia (NH3) was removed from the system. The stress is not enough
ammonia, so the equilibrium will shift to the right to replenish the ammonia removed.
Suppose some hydrogen (H2) was removed from the system. The stress is not enough
hydrogen, so the equilibrium will shift to the left to replenish the hydrogen removed.
Suppose the pressure is increased. The stress is too much pressure. In this case, the only
way to relieve the stress is to reduce the pressure. How may this be done? The pressure in the
container is due to the number of gas particles hitting the sides of the container with given force.
A reduction in the number of particles will cause the pressure to drop. Upon examining the
equilibrium, one sees that there are five moles of gas particles on the left hand side of the
equilibrium, while there are two moles of gas particles on the right hand side of the equilibrium.
Thus if the equilibrium shifts to the right, the number of gas particles in the container become
reduced, and the excess pressure is relieved.
Suppose a catalyst is added to the system. Since the catalyst is not a part of the
equilibrium, the equilibrium will not shift. The only affect that a catalyst has on an equilibrium is
to allow the system to reach equilibrium faster.
Suppose an inert gas, such as helium, as added to the system. Since helium is not a part of
this equilibrium, it will not affect the equilibrium.
It is important to note that the value of the equilibrium constant has not changed when the above
stresses were placed on the equilibrium. It is, after all, a constant. The only factor that will affect
the value of the equilibrium constant is temperature. In fact, you will note that when an
equilibrium constant is cited, the temperature at which that constant holds true is also cited. Thus
if the temperature is changed, the equilibrium shift will be accompanied by a change in the value
of the equilibrium constant.
Suppose you know that the following equilibrium is exothermic (
H2 (g) + Cl2 (g)
H < 0):
2 HCl (g) Keq = 4 x 1031 at 300 K
Since the process is exothermic, heat is released by the system to the surroundings, so one may
include heat in the equilibrium as a product:
H2 (g) + Cl2 (g)  2 HCl (g) + heat
Suppose the temperature of the system is increased to 500 K. The stress is too much heat,
so the equilibrium will shift to the left to remove the excess heat. In addition, the value of the
equilibrium constant decreases to 4 x 1018 , indicating the equilibrium lies further to the left.
When writing the mass action expression of an equilibrium, only include the components
in the equilibrium which are in the same phase. Never include solids. Let's look at some
equilibrium reactions involving more than one phase, and their mass action expressions.
CaCO3 (s)  CaO(s) + CO2(g)
Keq = [CO2]
Notice that the solids were not included as part of the mass action expression.
2 HCl(aq) + CaCO3 (s)
CaCl2 (aq) + CO2 (g) + 2 H2O(l)
The net ionic equation for the above reaction is:
2 H+(aq) + CaCO3(s)  Ca2+(aq) + CO2(g) + 2 H2O(l)
(notice that chloride ion is a spectator ion)
Keq1 
[Ca2  ]
, alternativ ely Keq2  [CO2]
[H ]2
Notice that the equilibrium mass action expression may be expressed in two ways; Keq1 is
expressed in terms of the concentrations of the aqueous species, while Keq2 is expressed
in terms of the concentration of the gaseous species. These two constants will have
different values. Which expression do you use? That depends on the information given in
the problem. If the concentrations of the aqueous species are given, use the expression for
Keq1 . If the concentration for CO2 is given, use the expression for Keq2 . Realistically, one
uses the expression whose quantities are most easily measured experimentally. The main
idea is not to mix phases, and that the solid is not included in either expression.
Quantitative Problems
Next, let's consider some quantitative problems dealing with chemical equilibria.
By examining the amounts of reactants and products at equilibrium, one may numerically
see how an equilibrium favors either reactants or products.
Suppose we are given the following equilibrium at 500 K:
CO(g) + 2 H2 (g)  CH3OH(g)
The equilibrium concentrations are: [CO] = 0.0911 M, [H2] = 0.0822 M,
[CH3OH] = 0.00892 M, what is the value of the equilibrium constant? Does the
equilibrium favor reactants or products?
First, we need to write the mass action expression:
Keq 
[CH3OH]
[CO][H2]2
Next, substitute the equilibrium concentrations into the mass action expression, and
calculate for the equilibrium constant:
Keq 
(0.00982 M)
2
(0.0911 M)(0.0822 M)
Since the value of the equilibrium constant is greater than one, (Keq >1), the equilibrium
favors the products.
Another type of equilibrium problem deals with finding the equilibrium
concentrations given the initial concentrations and the value of the equilibrium constant.
Suppose you are given the following equilibrium:
CO(g) + H2 O(g)  CO2 (g) + H2 (g)
Keq = 23.2 at 600 K
If the initial amounts of CO and H2O were both 0.100 M, what will be the amounts of
each reactant and product at equilibrium?
For this type of problem, it is convenient to set a table showing the initial
conditions, the change that has to take place to establish equilibrium and the final
equilibrium conditions. Let's begin by showing the initial conditions:
CO(g)
Initial
+
0.100
H2 O(g)

0.100
CO2 (g)
+
0
H2 (g)
0
Change
Equilibrium
Initially, 0.100 M CO and 0.100 M H2O are present. Equilibrium hasn't been established
yet, so the amounts of CO2 and H2 are assumed to be zero.
To establish equilibrium, some CO and H2O has to react, so we will call the amount of
CO and H2O reacted x, and the same x amount of CO2 and H2 must form:
CO(g)
Initial
Change
+
H2 O(g)

CO2 (g)
+
H2 (g)
0.100
0.100
0
0
-x
-x
+x
+x
Equilibrium
The amounts of reactants and products present at equilibrium will be the combination of
the initial amounts and the change. Just add the quantities together:
CO(g)
Initial
Change
+
H2 O(g)

CO2 (g)
+
H2 (g)
0.100
0.100
0
0
-x
-x
+x
+x
0.100 – x
Equilibrium
0.100 – x
x
x
Substitute the above algebraic quantities into the mass action expression:
Keq 
[CO2][H2]
(x)(x)
x2


 23.2
[CO][H2O] (0.100  x)(0.100  x)
(0.100  x}2
Since the algebraic expression is a perfect square, begin solving for x by taking the
square root of both sides of the equation:
x2

(0.100  x )2
23.2
x
 0.4.85
0.100  x )
Multiply both sides by the denominator, 0.100 - x:
(0.100  x ) x
x
 4.85((0.100  x )
(0.100  x )
The term, 0.100 - x, cancels out on the left hand side, and the term, 4.85 must be
distributed through the term 0.100 - x on the right hand side:
x = 0.485 – 4.85x
Add 4.85x to both sides of the equation:
x + 4.85x = 0.485 – 4.85x + 4.85 x
Combining terms:
5.85x = 0.485
Solve for x by dividing both sides by 5.85:
x 
0.485
 0.0829  [CO2]  [H2]
5.85
Recall that x represents the equilibrium quantities of both H2 and CO2 . The equilibrium
quantities of CO and H2O is given by:
0.100 – x = 0.100 – 0.0829 = 0.017 M = [CO] = [H2O]
Problems
1. Write the equilibrium constant expressions, Kc, for the following equilibria:
a.
N2(g) + 2 H2(g)  N2H4(g)
b. PF5(g)  PF3(g) + F2(g)
c. SiCl4(g) + 2 H2O(g)  SiO2(s) + 2 Cl2(g)
d. Ba3(PO4)2 (s)  3 Ba2+(aq) + 2 PO42-(aq)
Next write the Kp equilibrium constant expressions for the above equilibria, where applicable.
2. The following equilibrium has an equilibrium constant Kc = 1.7 x 10-3 at 2300 K and has a ΔH =
180.5 kJ/mol.
N2(g) + O2(g)
 2 NO(g)
a. Write the equilibrium constant expression in terms of concentrations and in terms of partial pressures.
b. Is this equilibrium reactant favored or product favored? Explain.
c. What is Kp for this reaction?
d. If [N2] = 0.10 M, [O2] = 3.2 M, and [NO] = 2.5 M are initially placed into the container, is the system
at equilibrium? If not, which way will the reaction have to shift to reach equilibrium?
e. The following changes are introduced to the equilibrium. Explain why and which direction the
reaction will shift in terms of LeChatelier’s Principle. You may assume the container is rigid. (1) N2 gas
is removed from the container. (2) NO gas is removed from the container. (3) O2 gas is added to
container. (4) Argon gas is added to the container. (5) The temperature of the system is decreased to
1000K. (6) The volume of the container is increased.
f. Which of the changes in problem (e) would alter the equilibrium constant value?
3. A sample of HI gas is placed into a rigid container at 700 K. At this temperature, Kc = 0.025. The
equilibrium is shown below:
2 HI(g)  H2(g) + I2(g)
a. Find the equilibrium concentrations of HI, H2 and I2 if the initial concentration of HI is 0.20 M.
b. Find the equilibrium concentrations of HI, H2 and I2 if the initial concentration of H2 = 0.40 M and the
initial concentration of I2 = 0.25 M.
c. What can be done to make the reaction product favored? Hint: Think LeChatelier’s Principle.
4. A sample of PCl5 gas is placed into a rigid 1.00 L container at 1500 K. At this temperature, Kc = 1.4 x
10-5. The equilibrium is shown below:
PF5(g)  PF3(g) + F2(g)
a. Find Kp at the same temperature.
b. If the initial concentration of PF5 is 0.30 M, what are the equilibrium concentration of all species in the
container?
c. If 0.0010 moles of F2 is introduced into the container, what will be the new equilibrium concentrations
of all species in the container?
Answers
1. Equilibrium constant expressions are always written as showing a ratio of products over reactants.
Here is a general format for a generic equilibrium:
aA + bB  cC + dD
The equilibrium constant, Kc , has the ratio expressed as molar concentrations:
[A]a [B]b
[C]c [D]d
Kc =
The equilibrium constant, Kc , is called an analytical equilibrium constant and it is applied to an analytical
equilibrium constant expression. The analytical equilibrium constant expression will only include those
components in the equilibrium that are in the same phase. Solids are never included because one cannot
measure the concentration of a solid.
The equilibrium constant, Kp , is applied to gas phase equilibria. The equilibrium constant expression is
expressed as a ratio in terms of the partial pressures of the gaseous components. Applying this to the
same equilibrium above, assuming A, B, C and D are all gases:
pcC ∙ pdD
Kp =
paA ∙ pbB
Use these concepts to write the equilibrium constant expressions for the equilibria in this question:
a. N2(g) + 2 H2(g)  N2H4(g)
This is a homogeneous equilibrium, meaning all the components are in the same phase, in this
the gas phase. Thus, we can write both a Kc expression and a Kp expression:
Kc =
[N2 H4 ]
[N2 ] [H2 ]2
and
Kp =
case,
pN2 H4
pN2 ∙ p2H2
b. PF5(g)  PF3(g) + F2(g)
This is also a homogeneous equilibrium, thus we can write both a Kc expression and Kp
Kc =
[PF3 ] [F2 ]
[PF5 ]
and K p =
expression:
pPF3 ∙ pF2
pPF5
c. SiCl4(g) + 2 H2O(g)  SiO2(s) + 2 Cl2(g)
This is a heterogeneous equilibrium. Solids are never included in the equilibrium constant expressions, so
we will only include the gases in this equilibrium. Since gases are involved, we are able to write both a
Kc and Kp equilibrium constant expression:
[Cl2 ]2
[SiCl4 ] [H2 O]2
BaSO4(s)  Ba2+(aq) + SO42- (aq)
Kc =
1. d
and
Kp =
p2Cl2
pSiCl4 p2H2 O
This is also a heterogeneous equilibrium. Exclude the solid barium sulfate and only include the aqueous
Ba2+ ions and SO42- ions. Since there isn’t a gas phase component only a Kc expression can be written:
Kc = [Ba2+] [SO42-]
Notice the absence of a denominator term!
2. a. Refer to the explanations in problem 1 to help you with writing the Kc and Kp, if needed.
Kc =
Kp =
p2NO
pN2 pO2
[NO]2
= 1.7 x 10−3
[N2 ] [O2 ]
= (value to be determined in part c of this problem)
b. The value of Kc = 1.7 x 10-3 is a small number (it is < 1). Mathematically this means the
denominator term must be large than the numerator term. Thus, there are more reactants (found in the
denominator term) than products (found in the numerator term), so the reaction must be reactant favored.
 cC(g) + dD(g), the relationship between Kp and Kc
c. Using a general reaction: aA(g) + bB(g)
is derived as follows:
From PV = nRT:
P =
n
RT
V
Since the n/V term is expressed in moles/liter, it is effectively the molarity, M, of the gas, P = MRT.
We can now write two equivalent expressions:
Kp =
pcC ∙ pdD
=
paA ∙ pbB
[C]c (RT)c [D]d (RT)d
[A]a (RT)a [B]b (RT)b
Regroup the expression on the right and hopefully you will recognize the equilibrium constant generated:
pcC ∙ pdD
paA ∙ pbB
=
[C]c [D]d
(RT)c (RT)d
∙
[A]a [B]b
(RT)a (RT)b
This shows how Kc is actually embedded in this expression, so the expression can be simplified to:
Kp = Kc (RT)Δn
where Δn = moles of gaseous products – moles of gaseous reactants. For the equilibrium
N2(g) + O2(g) 
2 NO(g)
Δn = 2 moles of gaseous NO – (1 mole of gaseous N2 + 1 mole of gaseous O2) = 0
Kp = 1.7 x 10-3[(0.08206 L∙atm/mol∙K)(2300K)]0 = 1.7 x 10-3
For this for this problem, Kp = Kc because any expression raised to the zero power = 1.
2 d. Use the reaction quotient, Qc, to determine whether the system is at equilibrium or whether the
system has to shift towards reactants or products to reach equilibrium. The Qc is evaluated using the
equilibrium constant expression. The actual initial amount of each component is introduced into the
equilibrium constant expression. The resulting constant, Qc, is then compared to the actual equilibrium
constant value, Kc. If Qc = Kc, the system is at equilibrium. If Qc > Kc, the numerator term must be too
large, so the system lies too far towards products. In order to reach equilibrium, the system must shift
towards reactants. If Qc < Kc, the denominator term is too large, so the system lies too far towards
reactants. In order to reach equilibrium, the system must shift towards products.
In this problem, Kc = 1.7 x 10-3. Calculate Qc and compare the value of Qc to Kc:
Qc =
[NO]2
(2.5)2
=
= 19
[N2 ][O2 ]
(0.10)(3.2)
Qc > Kc
19 > 1.7 x 10-3
As one can see, the system is not at equilibrium. The system currently lies too far towards products. In
order to reach equilibrium, the system must shift to towards reactants to reduce the amount of excess
products (the numerator must decrease) and increase the amount of reactants (the denominator must
increase).
e. LeChatelier’s Principle states that when a stress is placed on an equilibrium, the equilibrium shifts in
the direction that removes the stress. Based on this statement, it is important to first identify the stress.
The stress could be too much of something or too little of something. Once the stress is identified, it will
be easy to decide the direction of the shift.
(1) N2 gas is removed from the container. The stress created is not enough nitrogen gas in the container.
The equilibrium will shift towards reactants to replenish the nitrogen gas removed.
(2) NO gas is removed from the container. The stress created is not enough NO gas in the container.
The equilibrium will shift toward products to replenish the NO gas removed.
(3) O2 gas is added to the container. The stress created is too much O2 gas in the container. The
equilibrium will shift towards the products to remove the excess O2 gas by making more product.
(4) Argon gas is added to the container. Argon gas is not a part of this equilibrium. The partial pressures
of oxygen gas, nitrogen gas and nitrogen monoxide gas are unchanged. Only the total pressure increases.
Since the equilibrium doesn’t involve argon gas, no stress has been placed on the equilibrium so the
equilibrium doesn’t need to shift.
(5) The temperature of the system is decreased to 1000K. The reaction is endothermic; ΔH = +180.5
kJ/mol. This means heat is absorbed by the system from the surroundings. One may consider heat as a
reactant:
heat + N2(g) + O2(g)  2 NO(g)
The stress created is too little heat. The system will shift towards reactants to replenish the heat removed
by the action of cooling the system.
(6) The volume of the container is increased. According to Boyle’s Law, pressure and volume are
inversely proportional. Thus, increasing the volume will decrease the pressure. The stress could be too
little pressure and the system would want to shift to the side of the reaction with more moles of gas to
increase the particle number to bring the pressure back to equilibrium pressure. In this case, however, the
moles of gas particles is the same on both sides of the reaction, so pressure does not affect the
equilibrium.
2. f. The term equilibrium “constant” means exactly what is implied. The value is a constant, but only
at the temperature at which it was measures. Notice that when an equilibrium constant is given, a
temperature is always associated with that constant. The origin of the equilibrium constant is from the
rate constants of the forward and reverse reactions. Reaction rates are temperature dependent. Thus
when components of the equilibrium are introduced or removed, the equilibrium will shift, but as long as
the temperature is unchanged, the value of the equilibrium constant does not change. Hence only in
problem (5) above, will the constant change. If the equilibrium shifts towards reactants, as it does in this
problem, the Kc, will also decrease. If the equilibrium shifts towards products, the equilibrium constant,
Kc, will increase.
3. a. We need a reference point to start the problem. In this case, the assumption we make is that only
the HI is present initially; none of the H2 or I2 has formed:
Initial
2 HI(g)  H2(g) + I2(g)
0.20 M
0M
0M
There currently is a stress on this equilibrium; there are no products. In order to reach equilibrium, the
system has to shift towards products to reach equilibrium. For each 2 moles of HI that dissociates, 1 mole
of H2and 1 mole of I2 will from. Since we don’t know the exact amount of dissociation, we use the
mathematical amounts using “x” to represent the amount that shifts. Thus, for HI, a 2x amount is lost
(hence, -2x) and an x amount of both H2 and I2 is generated:
Initial
Change
2 HI(g)  H2(g) + I2(g)
0.20 M
0M
0M
-2x
+x
+x
Combining the Initial and the changed amounts, we have the amount at equilibrium:
2 HI(g)
Initial
0.20 M
Change
-2x
Equilibrium 0.20 – x
 H2(g) + I2(g)
0M
0M
+x
+x
+x
+x
The equilibrium constant expression is written as:
Kc =
[H2 ][I2 ]
= 0.025
[HI]2
Substitute the algebraic expressions at equilibrium in to the equilibrium constant expression:
(x)(x)
= 0.025
(0.20 – x )2
Sometimes one can simplify the expression by assuming that the x amount found in the denominator term
is negligibly small. Compare the magnitude of the known concentration to the magnitude of the
equilibrium constant. If constant is smaller by a factor of 100 or more, the x being considered can be
neglected as being negligibly small. In this case, the difference between the magnitude of the
concentration and the equilibrium constant is less than 100, so x cannot be neglected, an a quadratic
equation must be solved. We have to covert the equilibrium constant to quadratic form. The series of
steps needed to do so is as follows:
Take the denominator term and turn it into a quadratic by performing the appropriate multiplications:
x2
x2
=
= 0.025
(0.20 – x)(0.20 – x)
0.040 – 0.40x + x 2
Multiply both sides of the equation by the denominator:
(0.040 − 0.40x + x 2 ) x
x2
= 0.025(0.040 − 0.40x + x 2 )
(0.040 − 0.40x + x 2 )
Distribute 0.025 through on the right hand side of the expression:
x 2 = 0.0010 − 0.010x + 0.025x 2
Bring all terms to the same side of the equation. In this case, x 2 is subtracted from both sides of the
expression:
0 = 0.0010 – 0.010x – 0.975x2
This expression as quadratic form: ax2 + bx + c = 0, where a = -0.975, b = – 0.01 and c = 0.0010.
Use the quadratic formula to solve for both solutions of x:
x± =
−b ± √b 2 − 4ac
2a
−(−0.010) ± √(−0.010)2 – 4(−0.975)(0.0010)
2(−0.975)
x± =
Work the arithmetic under the radical first:
x± =
0.010 ± √0.0040
0.010 ± 0.063
=
−1.95
1.95
Find the two solutions to x:
x+ =
0.010 + 0.063
= − 0.038
−1.95
x− =
0.010 – 0.063
= 0.027
−1.95
Only one of these two solutions makes any physical sense (x = 0.027). The other solution for x,
x = –0.038, is impossible, as it would give an impossible amount for the equilibrium concentration of HI.
One can see this by applying x to the equilibrium amounts:
For x = - 0.038, [HI] = 0.20 – 2x = 0.20 – 2(-0.038) = 0.28 M (more than one started with)
[H2] = [I2] = x = -0.038M (negative concentration values are impossible)
For x = 0.027, [HI] = 0.20 – 2x = 0.20 – 2(0.027) = 0.15 M
[H2] = [I2] = 0.027 M
(these values make “physical” sense)
An alternative method to solve this problem is to recognize that your algebraic expression is a
perfect square:
(x)(x)
= 0.025
(0.20 – x )2
Take the square root of both sides of the expression to simplify the expression:
(x)(x)
√
= √0.025
(0.20 – x )2
x
= 0.16
0.20 – x
The above expression is much easier to solve for x. Start by multiplying both sides by the
denominator term to have it cancel on the left hand side of the expression:
(0.20 – x) 
x
= 0.16(0.20 – x)
(0.20 – x)
Distribute the 0.16 through on the right hand side of the expression:
x = 0.032 – 0.16x
Add 0.16 x to both sides to collect the x terms:
x + 0.16x = 0.032 – 0.16x + 0.16x
1.16x = 0.032
Divide both sides by 1.16 to solve for x:
1.16x
0.032
=
1.16
1.16
x = 0.028 = [H2] = [I2]
0.20 -2x = 0.20 – 2(0.028) = 0.14 = [HI]
These values agree in the last decimal place, an acceptable error.
3. b. The solution to this problem will be analogous to the solution in the previous part of this
question. In this case one must use the quadratic equation to solve the problem since we do not
have a perfect square.
Set up the table as we did in part a of this problem:
Initial
2 HI(g)
0M
 H2(g) + I2(g)
0.40M 0.20 M
Notice that initial amount of HI is zero, and that we have initial amounts of “products” as our reference
point. This helps us decide which way the equilibrium has to shift to reach equilibrium. The stress is the
lack of HI, so an unknown amount of x has to be lost by H2 and I2 to form twice as much, 2x, of HI:
Initial
Change
2 HI(g)
0M
+2x
 H2(g) + I2(g)
0.40M
0.20 M
-x
-x
Initial
Change
Equilibrium
2 HI(g)
0M
+2x
2x
 H2(g) + I2(g)
0.40M
0.20 M
-x
-x
0.40 – x
0.20 – x
Thus at equilibrium:
The equilibrium constant expression is written as:
Kc =
[H2 ][I2 ]
= 0.025
[HI]2
Substitute the algebraic expressions at equilibrium in to the equilibrium constant expression:
(0.40 − x)(0.20 – x)
= 0.025
2x
Multiply out the factors in the numerator:
0.080 – 0.60x + x 2
= 0.025
2x
Multipy through by 2x to cancel the denominator term on the left hand side of the expression:
0.080 – 0.60x + x 2 2x

= 0.025(2x)
2x
1
0.080 – 0.60x + x2 = 0.050x
Subtract 0.050x from both sides of the expression to get the equation into quadratic form:
0.080 – 0.60x – 0.050x + x2 = 0.050x – 0.050x
The equation in quadratic form is 0.080 – 0.65x + x2 = 0, where a = 1, b = -0.65 and c = 0.080.
Substitute these values into the quadratic equation:
x± =
x± =
−b ± √b 2 − 4ac
2a
−( −0.65) ± √(−0.65)2 – 4(1)(0.080)
2(1)
Find the solution under the radical first then:
x± =
0.65 ± √0.1025
0.65 ± 0.32
=
2
2
Solve for the two possible values of x:
x+ =
0.65 + 0.32
= 0.48
2
x− =
0.65 – 0.32
= 0.16
2
Of these two solutions, 0.48 is impossible because it is greater than the original concentrations of H2
or I2. Thus, the only answer that makes physical sense is 0.16 M. Use this to find the equilibrium
amounts:
[H2] = 0.40 – 2x = 0.40 –0.16 = 0.24 M
[I2] = 0.20 – x = 0.20 – 0.16 = 0.04 M
[HI] = 2x = 2(0.16) = 0.32M
3. c. If the proper stress is placed on the equilibrium, the equilibrium could be made to shift
towards products. The stresses that could be used are: (1) Addition of HI, (2) removal of H2, or (3)
removal of I2. A change in volume would not help here, as the moles of gaseous reactants equals the
moles of gaseous products, so a pressure change as a result of a change in the volume of the
container would not work. Since a enthalpy value has not been given, one cannot predict how an
increase or decrease in temperature would impact the equilibrium position.
4. a. The relationship between Kp and Kc is: Kp = Kc(RT)Δn (see problem 2c for the derivation).
Our reaction is:
PF5(g)  PF3(g) + F2(g)
Kp = 1.4 x 10-5[(0.08206 L∙atm/mol∙K)(1500K)]1 = 1.7 x 10-3
where Δn for this reaction is:
2 moles of gaseous products – 1 mole of gaseous reactants = 1
4. b. Set up the table of initial conditions to provide a reference point:
PF5(g)  PF3(g) + F2(g)
Initial 0.30
0
0
The stress on this equilibrium at present is the lack of products; the reaction must shift towards
products by an x amount:
PF5(g)  PF3(g) + F2(g)
Initial 0.30
0
0
Change -x
+x
+x
The equilibrium amounts are obtained by combining the initial amount and the change:
PF5(g)  PF3(g) + F2(g)
Initial
0.30
0
0
Change
-x
+x
+x
Equilibrium 0.30 –x
x
x
The equilibrium constant expression is:
Kc =
[PF3 ][F2 ]
[PF5 ]
Substitute the algebraic expressions into the equilibrium constant expression:
(x)(x)
= 1.4 x 10−5
0.30 − x
To solve for x, one can take advantage of the fact the equilibrium constant is small. This means that
the x amount of PF5 lost will be negligibly small and can be ignored. The expression simplifies to:
x2
= 1.4 x 10−5
0.30
Solve for x by first multiplying through by the denominator:
0.30
x2
∙
= (0.30)(4.2 10−6 )
1
0.30
√x 2 = √4.2 10−6
Next take the square root of both sides of the equation to solve for x. The value of x represents the
equilibrium amounts of PF3 and F2. The amount of PF5 is 0.30 –x, and one can see that it was
negligibly small:
x = 2.0 x 10-3M = [PF3] = [F2]
[PF5] = 0.30 – 0.002 = 0.30M
4. c. Since the container is 1.0 L in size, the new initial concentration of F2 will be 0.0010M. Set up
the table showing the initial conditions , change and equilibrium conditions as before:
PF5(g)  PF3(g) + F2(g)
Initial
0.30
0
0.0010
Change
-x
+x
+x
Equilibrium 0.30 –x
x
0.0010 + x
The equilibrium constant expression is:
Kc =
[PF3 ][F2 ]
[PF5 ]
Substitute the algebraic expressions into the equilibrium constant expression:
(x)(0.0010 + x)
= 1.4 x 10−5
0.30 − x
The x amount of additional F2 in the expression, 0.0010 + x, is not negligibly small. Although the x
in the denominator is negligible, and since we have to use the quadratic equation, we might as well
not neglect any of the variables. Multiply both sides by the denominator to start converting the
expression to quadratic form:
(0.30 − x) (x)(0.0010 + x)
(0.30 – x)
∙
= 1.4 x 10−5 ∙
1
0.30 − x
1
Distribute terms where needed:
0.0010x + x2 = 4.2 x 10-6 – (1.4 x 10-5)x
Collect terms and bring them to same side of the equation:
x2 + 0.0010x + (1.4 x 10-5)x – 4. 2 x 10-6 = 4.2 x 10-6 – (1.4 x 10-5)x – 4.2 x 10-6 + (1.4 x 10-5)x
The above equation simplies to quadratic form:
x2 + 1.014 x 10-3x – 4.2 x 10-6 = 0
where a = 1, b = 1.014 x 10-3 and c = – 4.2 x 10-6
Substitute these values into the quadratic equation:
x± =
−b ± √b 2 − 4ac
2a
−( 1.014 x 10−3 ) ± √(1.014 x 10−3 )2 – 4(1)(−4.2 x 10−6 )
2(1)
Find the solution under the radical first then:
x± =
x± =
−( 1.014 x 10−3 ) ± √1.78 x 10−5
− 1.014 x 10−3 ± 4.2 x 10−3
=
2(1)
2(1)
Solve for the two possible values of x:
x+ =
− 1.014 x 10−3 + 4.2 x 10−3
= 0.0016
2
x− =
− 1.014 x 10−3 − 4.2 x 10−3
= −0.0026
2
A concentration cannot be a negative value, so –0.0026 is an impossible result. Thus, the
equilibrium concentrations of all species are:
[PF5] = 0.30 – x = 0.30 – 0.0016 = 0.30M
[F2] = 0.0010 + x = 0.0010 + 0.0016 = 0.0026M
[PF3] = x = 0.0016 M