Chemical Equilibrium
Equilibrium
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A state where the reactants and products remain
constant over time.
For some reactions, the equilibrium position
favors the products and the reaction appears to
have gone to completion (amount of reactants is
negligible)-”Equilibrium lies to the right” (in
direction of the products)
Example: 2 H2 + O2  2 H2O
Other reactions only occur to a small extent with
the product virtually undectable-”Equilibrium lies
to the left” (in direction of the reactants)
Example: 2 CaO  2Ca + O2
Equilibrium is not static


Because the concentrations do not
change, it appears that the reaction has
stopped.
Instead, equilibrium is highly dynamic with
the reactions continuing to occur in both
directions at the same rate.
products
reactants
As the
concentration of
the reactants
decreases, the
forward reaction
slows down.
As the
concentration of
the products
increases, the rate
of the reverse
reaction increases
Eventually, the
rates become
equal.
Factors affecting equilibrium
position:
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Initial concentrations***
Energies of reactants and products
“organization” of reactants and products.
***The factor to be addressed in this unit.
Equilibrium Constants
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Through experimentation and observation, the
Law of Mass Action was proposed.
The law suggests that for a reaction of the
following type:
jA + kB  lC + mD
the following equilibrium expression is used to
represent the reaction.
K=
[C]l
[D]m
[A]j [B]k
K is the equilibrium constant
[ ] represent the
concentrations at equilibrium
Coefficients become the
exponents
Practice

Write the equilibrium expression for the
following equation:
4 NH3 + 7 O2 < -- > 4 NO2 + 6 H2O
Determining the Equilibrium
Constant
The equilibrium constant can be calculated at a
given temperature if the equilibrium
concentrations of the reaction components are
known.
N2 + 3H2 < -- > 2 NH3
[NH3] = 3.1 x 10-2 M
[N2] = 8.5 x 10-1 M
[H2] = 3.1 x 10-3 M

Calculate the value for K
Other methods for Determining
Equilibrium Constants

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If the reaction is reversed,
2NH3 < --> N2 + 3H2
then, K/ = 1/K
If the reaction is multiplied by a factor,
½ N2 + 3/2 H2 < -- > NH3
then, K// = (K)1/2
Application from the lab
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At a given temperature:
1) K always has the same value regardless of the
starting concentrations
2) the equilibrium concentrations will not always be the
same
(The set of equilibrium concentrations is called the
equilibrium position.)
There is only one equilibrium constant at a particular
temperature, but there is an infinite number of
equilibrium positions.
The specific equilibrium position depends on the starting
concentrations, but the equilibrium constant does not.
See example 13.3
Heterogeneous Equilibrium
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Many reactions involve reaction components in
more than one phase.
Example: CaCO3 (s) < -- > CaO (s) +CO2 (g)
Because the concentrations of pure solids and
liquids cannot change, they are not included in
the equilibrium expression.
K = [CO2]
Write the equilibrium expression for the
following equation:
PCl5(s) < -- > PCl3(l) + Cl2(g)
Homework


The following problems may now be
completed:
17, 20, 22, 24, and 29 on pages 614-615
Equilibrium Involving Pressure
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

Equilibrium involving gases can be
described in terms of pressures (as well as
concentrations)
PV = nRT or P = (n/V)RT where n/V is
the concentration of the gas.
The equilibrium constant in terms of
partial pressures in written as KP
Sample Problem

Given the following reaction:
2NO(g) + Cl2 (g) < -- > 2NOCl(g)
PNOCl = 1.2 atm
PNO = 0.050 atm
PCl2 = 0.30 atm
Calculate the value of Kp
Relationship between K and KP
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KP = K(RT)Δn
(see pages 587-588 to view how this equation
was derived)
R = 0.08206
T is the temperature in Kelvin
Δn = the difference in the sum of the
coefficients for the products and the reactants.
Example: jA + kB < -- > lC + mD
Δn = (l + m) – (j + k)
Sample Problem

Using the value for Kp obtained in the
previous sample problem, calculate the
value of K for the reaction
2NO(g) + Cl2(g) < -- > 2NOCl(g)
Applications of the Equilibrium
Constant

Knowing the equilibrium constant allows us to
predict several important features of the
reaction.
1) the tendency of the reaction to occur (but not
the speed)
2) whether a given set of concentrations
represent an equilibrium condition
3) the equilibrium position that will be achieved
from a given set of initial concentrations.
The Extent of a Reaction
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The tendency for a reaction to occur is indicated
by the magnitude of the equilibrium constant.
A value of K much larger than 1 means that at
equilibrium the reaction will consist mostly of
product-the equilibrium lies to the right. (The
reaction goes to completion)
A value of K much smaller than 1 means that at
equilibrium the reaction will consist mostly of
reactants-the equilibrium lies to the left. (The
reaction does not occur to any given extent)
The size of K and the time required to reach
equilibrium are not directly related.
Reaction Quotient
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When given a set of reaction components, it is helpful to
know if the mixture is at equilibrium or, if not, in what
direction the system must shift to reach equilibrium.
To determine the direction of the move toward
equilibrium, we use the reaction quotient (Q).
The reaction quotient is obtained by applying the law of
mass action to the initial concentrations instead of
equilibrium concentrations.
If Q = K, the system is at equilibrium; no shift will occur
If Q > K, the system shifts to the left.
If Q < K, the system shifts to the right.
Complete the sample problem on page 593.
Le Chatelier’s Principle

Le Chatelier’s Principle states that if a
change is imposed on a system at
equilibrium, the position of equilibrium will
shift in a direction that tends to reduce
that change.
Effect of a Change in
Concentration
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
If a component is added to a reaction
system at equilibrium (at constant T and
P), the equilibrium position will shift in the
direction that lowers the concentration of
that component (away from).
If a component is removed from the
system, the system will shift in the
direction that increases the concentration
of that component (towards)
Copy the following equation:

As4O6(s) + 6C(s) < -- > As4(g) + 6CO (g)
In which direction will the equilibrium
position shift if CO is added?
0%
1.
0%
2.
0%
3.
Left
Right
No shift will occur
1
2
3
4
5
6
7
8
9
10
21
22
23
24
25
26
27
28
29
30
11
12
10
13
14
15
16
17
18
19
20
In which direction will the equilibrium
position shift if C is removed?
0%
1.
0%
2.
0%
3.
Left
Right
No shift will occur
1
2
3
4
5
6
7
8
9
10
21
22
23
24
25
26
27
28
29
30
11
12
10
13
14
15
16
17
18
19
20
In which direction will the equilibrium
position shift if CO is removed?
0%
1.
0%
2.
0%
3.
Left
Right
No shift will occur
1
2
3
4
5
6
7
8
9
10
21
22
23
24
25
26
27
28
29
30
11
12
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13
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16
17
18
19
20
Effect of a Change in Pressure

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Adding an inert gas has no effect on the
equilibrium position (has no effect on the
concentrations or partial pressures of the
reactants or products).
Increasing the volume of the container
decreases the pressure. The system responds
by increasing the pressure through the
production of more gaseous molecules (a shift
toward the side with the greatest number of gas
molecules).
Decreasing the volume of the container results
in the opposite occurrence.
What shift in the equilibrium position will occur if
the volume is reduced for the following process:
P4(s) + 6Cl2(g) < -- > 4PCl3(l)
0%
1.
0%
2.
0%
3.
Left
Right
No shift will occur
1
2
3
4
5
6
7
8
9
10
21
22
23
24
25
26
27
28
29
30
11
12
10
13
14
15
16
17
18
19
20
What shift in the equilibrium position will occur if
the volume is reduced for the following process:
PCl3 (g) + Cl2 (g) < -- > PCl5 (g)
0%
1.
0%
2.
0%
3.
Left
Right
No shift will occur
1
2
3
4
5
6
7
8
9
10
21
22
23
24
25
26
27
28
29
30
11
12
10
13
14
15
16
17
18
19
20
What shift in the equilibrium position will occur if
the volume is reduced for the following process:
PCl3(g) + 3NH3(g) < -- > P(NH2)3(g) + 3HCl(g)
0%
1.
0%
2.
0%
3.
Left
Right
No shift will occur
1
2
3
4
5
6
7
8
9
10
21
22
23
24
25
26
27
28
29
30
11
12
10
13
14
15
16
17
18
19
20
Effect of a Change in Temperature

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Changing the temperature changes the
value of K.
In exothermic reactions (heat is given off
and is therefore written as a product), the
equilibrium position will shift to the left if
the temperature is increased and to the
right if the temperature is decreased.
For endothermic reactions, the opposite
will occur.
For the following reaction, predict how the value of
K changes as the temperature is increased.
N2(g) + O2(g) < -- > 2NO(g) ΔH = 181 kJ
0%
1.
0%
2.
0%
3.
Increases
Decreases
Remains the same
1
2
3
4
5
6
7
8
9
10
21
22
23
24
25
26
27
28
29
30
11
12
10
13
14
15
16
17
18
19
20
For the following reaction, predict how the value of
K changes as the temperature in increased.
2SO2(g) + O2(g) < -- > 2SO3(g) ΔH = -198 kJ
0%
1.
0%
2.
0%
3.
Increases
Decreases
Remains the same
1
2
3
4
5
6
7
8
9
10
21
22
23
24
25
26
27
28
29
30
11
12
10
13
14
15
16
17
18
19
20
Copy the following equation

N2(g) + 3H2(g) < -- > 2NH3(g) + 92.94 kJ
How will the equilibrium position shift if the
temperature increases?
0%
1.
0%
2.
0%
3.
Left
Right
No shift will occur
1
2
3
4
5
6
7
8
9
10
21
22
23
24
25
26
27
28
29
30
11
12
10
13
14
15
16
17
18
19
20
How will the equilibrium position
shift if the volume increases?
0%
1.
0%
2.
0%
3.
Left
Right
No shift will occur
1
2
3
4
5
6
7
8
9
10
21
22
23
24
25
26
27
28
29
30
11
12
10
13
14
15
16
17
18
19
20
How will the equilibrium position
shift if NH3 is removed?
0%
1.
0%
2.
0%
3.
Left
Right
No shift will occur
1
2
3
4
5
6
7
8
9
10
21
22
23
24
25
26
27
28
29
30
11
12
10
13
14
15
16
17
18
19
20
How will the equilibrium position
shift if N2 is added?
0%
1.
0%
2.
0%
3.
Left
Right
No shift will occur
1
2
3
4
5
6
7
8
9
10
21
22
23
24
25
26
27
28
29
30
11
12
10
13
14
15
16
17
18
19
20
How will the equilibrium position
shift if some Ar(g) is added?
0%
1.
0%
2.
0%
3.
Left
Right
No shift will occur
1
2
3
4
5
6
7
8
9
10
21
22
23
24
25
26
27
28
29
30
11
12
10
13
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15
16
17
18
19
20
Homework

All homework problems that accompany
Chapter 13 should be completed for class
tomorrow.