AP Chapter 13 Notes
Chemical Equilibrium
Many reactions do not go to completion. They will only go so far and
then the concentration of the products remains
.
Chemical equilibrium is the state where …..
….in fact, any reaction will reach equilibrium if left in a closed container
over time.
The equilibrium position may favor products so the reaction appears…..
Or the equilibrium position may favor the reactants so that very little
are formed.
A reaction system is at equilibrium when:
We can see a reaction equation written at equilibrium as:
H2O + CO  H2 + CO2
What happens if more H2O is added?
First:
But then:
Equilibrium Position is affected by:
 Initial
 Relative energies of
 Relative degree of
 (nature wants to
aA + bB  eE + fF
Q
=
.
.
.
)
can be written in a mass action expression:
.
If the reaction is at equilibrium, we write the expression as:
K is the equilibrium constant and has no units.
Example 13.1
Write the equilibrium expression for the following reaction:
4NH3 (g) + 7O2 (g)  4NO2 (g) + 6H2O (g)
Example 13.2
The following equilibrium concentrations were observed for the Haber
process at 127oC:
[NH3] = 3.1 x 10-2 mol/L
[N2]= 8.5 x 10-1 mol/L
[H2] = 3.1 x 10-3 mol/L
a. calculate the value of K at 127oC for N2 (g) + 3H2 (g)  2NH3 (g)
b. Calculate the value for K at 127oC for the reaction:
2NH3 (g)  N2 (g) + 3H2 (g)
c. calculate K at 127oC for the reaction given:
½ N2 (g) + 3/4 H2 (g)  NH3 (g)
With the original reaction:
aA + bB  eE
+ fF
…the equilibrium expression is:
…and the reverse reaction is:
…or 1/K
This just means that at a particular temperature, the ratio of the product
of the products over the product of the reactants is a constant, K. The
may vary, but the constant remains the same.
Example 13.3
The following results were collected for two experiments involving the
reaction at 600 oC between gaseous sulfur dioxide and oxygen to form
gaseous sulfur trioxide:
Experiment 1
Experiment 2
Initial
Equilibrium
Initial
Equilibrium
[SO2]o = 2.00 M
[SO2] = 1.50 M
[SO2]o = 0.500 M
[SO2] = 0.590 M
[O2]o = 1.50 M
[O2] = 1.25 M
[O2]o = 0 M
[O2] = 0.0450 M
[SO3]o = 3.00 M
[SO3] = 3.50 M
[SO3]o = 0.350 M
[SO3] = 0.260 M
Show that the equilibrium constant is the same in both cases.
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Q is the
expression
are represents the mass action
.
Q = Kc only at
If Q < Kc
If Q > Kc
Equilibrium involving gases can be expressed in terms of pressure:
Remember that
PV = nRT
or manipulating we get:
P = n RT = CRT
V
….where C is the factor n/V
…so for
N2 + 3H2 
or molar concentration
2NH3
Kp =
=
=
.
Example 13.4
The reaction for the formation of nitrosyl chloride 2NO (g) + Cl2 (g)
2NOCl (g) was studied at 25 oC. The pressures at equilibrium were
founded to be PNOCl = 1.2 atm PNO = 5.0 x 10-2 atm PCl2 = 3.0 x 10-1
atm. Calculate the value of Kp for this reaction at 25oC.
Kc & Kp
Kc and Kp do not have the same value and do not occur at the same point.
N2 + 3H2  2NH3
Kc =
3
.
Kp =
.
The equation equating the two is:
Kp = Kc (RT)n
where n is
.
…so for the above reaction, n =
What is the Kc for the reaction PCl3(g) + Cl2(g)  PCl5(g) at 25oC when
the Kp= 5.6 x 106 ? (R = 0.0821 Latm/(molK)
Heterogeneous Equilibria
For heterogeneous reactions, the equilibrium expression does not include
pure substances. Why not?
What are the values for kp and kc for the reaction: C6H6(l)  C6H6(g)
at 25oC, given that the vapor pressure of benzene at 25oC equals 92 torr?
Example 13.6
Write the expressions for K and Kp for the following:
a. The decomposition of solid phosphorus pentachloride to liquid
phosphorus trichloride and chlorine gas
b. deep blue solid copper(II) sulfate pentahydrate is heated to drive off
water vapor to form white solid copper(II)sulfate.
Knowing K, we can predict:

of a reaction to occur.
 If concentrations represent
.

that will be achieved from
initial concentrations.
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Tendency:
We can tell a lot by the magnitude of K….
Large K:
Small K:
Concentrations:
What happens if we’re only given initial concentrations and need
to find equilibrium concentrations?
Example
If 0.040 mole of H2 and 0.025 mole of I2 are placed in a one-liter container
and allowed to come to equilibrium, the equilibrium concentration of I2 is
0.0022M. Calculate the equilibrium concentration of H2 and HI. Find Kc.
H 2(g) + I 2(g)  2 HI (g)
Initial [ ]
 [ ]
final [ ]
.
[H2]=
[HI]=
Kc =
Example
Some pure phosgene is placed in a closed reaction container. At
equilibrium, the partial pressure of phosgene in the container is 0.25 atm.
What are the partial pressures of CO(g) and Cl2(g)?
Kp = 4.6x10 -2 atm at 395oC.
COCl2(g)  CO(g) + Cl2(g)
Initial [ ]
 [ ]
final [ ]
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.
Example
The equilibrium constant at 425oC is 53. If 0.040 moles each of H2 and I2
are placed in a one liter flask with 0.75 moles of HI at 425oC, is the system
at equilibrium? Will the forward or reverse reaction occur to greater
extent to establish equilibrium?
H2(g) + I2(g)  2 HI(g)
Initial [ ]
 [ ]
.
final [ ]
Example
The equilibrium constant at 425oC is 53. If 0.150 moles each of H2 and I2
are placed in a 5.00 liter flask and heated to 425oC and allowed to reach
equilibrium, what will be the concentration of each substance at
equilibrium?
H2(g) + I2(g)  2 HI(g)
Initial [ ]
 [ ]
.
final [ ]
[H2]=
[I2]=
[HI]=
Example
The equilibrium constant at 425oC is 53. If 0.300 moles of HI is placed in
a 5.00 liter flask and heated to 425oC and allowed to reach equilibrium,
what will be the concentration of each substance at equilibrium?
H2(g) + I2(g)  2 HI(g)
Initial [ ]
 [ ]
.
final [ ]
[H2]=
[I2]=
[HI]=
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Example
The equilibrium constant at 300oC is 2.0x10-26. If 0.50 moles of SO3 is
placed in a 10.00 liter flask, what will be the equilibrium concentration of
the three gases at 300 K?
2SO3(g)  2 SO2(g) + O2(g)
Initial [ ]
 [ ]
.
final [ ]
[SO3]=
[SO2]=
[O2]=
check the 5% rule
Le Chatelier’s Principle!
If a system in dynamic equilibrium is subject to a disturbance that upsets
the equilibrium, the system responds in such a way as to
the
disturbance and if possible return to equilibrium.
2 SO2(g) + O2(g)  2 SO3(g) + 188 kJ
If more SO2 added
If more SO3 added,
If O2 taken out,
If heat added,
If cooled,
N2O4(g)  2 NO2(g)
colorless
reddish brown
The color becomes darker red with an increase in temperature. Is the
reaction exothermic or endothermic?
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Predict the direction in which the system will shift to reach equilibrium in
each of the following:
Example 13.7
For the synthesis of ammonia at 500 oC, the equilibrium constant is
6.0 x 10-2.
N2 + 3H2  2NH3
[NH3]=1.0x10-3
[N2]=1.0x10-5
[H2]=2.0x10-3
Example 13.8
Gaseous N2O4 was placed in a flask and allowed to reach equilibrium at a
temperature where Kp =0.133. At equilibrium, the pressure of N2O4 was
found to be 2.71 atm. Calculate the equilibrium pressure of NO2(g).
N2O4(g) 
2NO2(g)
Example 13.9
At a certain temperature a 1.00 L flask initially contained 0.298 mol
PCl3(g) and 8.70 x 10-3 mol PCl5 (g). After the system had reached
equilibrium, 2.00x10-3 mol Cl2 (g) was found in the flask. Calculate the
equilibrium concentrations of all species and the value of K.
Pressure and volume changes do affect
.
How does changing the volume of a container of gas affect the
equilibrium?
Does adding an inert gas affect the equilibrium of a reaction? Why or why
not?
What effect does adding a catalyst have on a reaction at equilibrium?
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Additional Examples not on presentation:
Example 13.10
Carbon monoxide reacts with steam to produce carbon dioxide and
hydrogen. At 700 K the equilibrium constant is 5.10. Calculate the
equilibrium concentrations of all species if 1.000 mol of each component
is mixed in a 1.000-L flask.
Example 13.11
Assume that the reaction for the formation of gaseous hydrogen fluoride
from hydrogen and fluorine has an equilibrium constant of 1.15 x 102 at a
certain temperature. In a particular experiment, 3.000 mol of each
component was added to a 1.500-L flask. Calculate the equilibrium
concentrations of all species.
Example 13.12
Assume that gaseous hydrogen iodide is synthesized from hydrogen gas
and iodine vapor at a temperature where the equilibrium constant is
1.00x102. Suppose HI at 5.000x 0-1 atm, H2 at 1.000x10-2 atm, and I2 at
5.000x10-3 atm are mixed in a 5.000-L flask. Calculate the equilibrium
pressures of all species.
Example 13.13
Arsenic can be extracted from its ores by first reacting the ore with oxygen
(called roasting) to form solid As4O6, which is then reduced using carbon:
As4O6 (s) + 6C (s) 
As4 (g) + 6CO (g)
Predict the direction of the shift of the equilibrium position in response to
each of the following changes in conditions:
a. Addition of carbon monoxide
b. Addition or removal of carbon or tetraarsenic hexoxide (As4O6)
c. Removal of gaseous arsenic (As4)
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Example 13.14
Predict the shift in equilibrium position that will occur for each of the
following processes when the volume is reduced.
a. The preparation of liquid phosphorus trichloride by the reaction
P4 (s) + 6Cl2 (g) 
4PCl3 (l)
b. The preparation of gaseous phosphorus pentachloride according to the
equation:
PCl3 (g) + Cl2 (g) 
PCl5 (g)
c. The reaction of phosphorus trichloride with ammonia:
PCl3 (g) + 3NH3 (g) 
P(NH2)3 (g) + 3HCl (g)
Example 13.15
For each of the following reactions, predict how the value of K changes as
the temperature is increased.
a. N2 (g) + O2 (g) 
2NO (g)
 Ho = 181 kJ
b. 2SO2 (g) + O2 (g)

2SO3 (g)
 Ho = -198 kJ
Assignment Chapter 13:
# 9, 19, 22, 23, 26, 28, 31, 34, 37, 41, 44, 46, 50, 53, 56, 60, 63
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