Ch 14 Chemical Equilibrium
Equilibrium
 A chemical reaction that “stops”
before it is completed
 Mixture of reactant and product
 Reaction is “reversible” and noted
with a double arrow
 Means that the forward reaction and
reverse reaction occur at the same
rate…DYNAMIC EQUILIBRIUM
Reactions are reversible
A + B
C + D ( forward)
C + D
A + B (reverse)
 Overall A + B
C+D
 Initially there is only A and B so only
the forward reaction is possible
 As C and D build up, the reverse
reaction speeds up while the forward
reaction slows down.
 Eventually the rates are equal
1
Reaction Rate
Forward Reaction
Equilibrium
Reverse reaction
Time
What is equal at Equilibrium?
 Rates are equal
 Concentrations are not.
 Rates are determined by
concentrations and activation energy.
 The concentrations do not change at
equilibrium, definite concentrations.
 Time it takes to reach equilibrium is
not a factor.
Example
 CO(g) + 3H2(g)
CH4(g) + H2O(g)
 At time t=0 sec
 Given 1.000 moles carbon monoxide and 3.000 moles
hydrogen gas in a 10.00 L vessel @ 1200 K
 How much of each product do you have?
 What’s going to happen?
2
Ultimately, what do we
want to know about a reaction
that reaches equilibrium?
 HOW MUCH OF EACH
SUBSTANCE IS IN THE
CONTAINER AT EQUILIBRIUM?
Law of Mass Action
 is a mathematical model that explains
and predicts behaviors of substances
in dynamic equilibrium.
 For a reaction at a given temperature,
the behavior is the same all the time,
its constant
 We call it EQUILIBRIUM CONSTANT
 We symbolize it with a Keq
Law of Mass Action
 For any reaction
 jA + kB
 K = [C]l[D]m
[A]j[B]k
lC + mD
PRODUCTSpower
REACTANTSpower
 This says that a reaction, at the given
temperature, will reach equilibrium, and
the ratio of the concentration of the
products to the reactants will ALWAYS be
the same
3
Equlibrium Constant, K
jA + kB
lC + mD
l
K = [C] [D]m
[A]j[B]k
where [ ] means concentration as molarity of
substances A, B, C, D at equilibrium
ONLY involves substances that change concentration
which includes solutions (aq) and gases (g)
Therefore, pure solids & liquids are not involved in the
equation
Write the equilibrium constant expression
What is K if rxn is reversed?
 Original Reaction
jA + kB
lC + mD
 If we write the reaction in reverse.
lC + mD
jA + kB
 Then the new equilibrium constant is
K’ = [A]j[B]k = 1/K
[C]l[D]m
4
What is the value of K if…
Given N2O4(g)2 NO2(g)
Kc = 0.212 @ 100oC
What is the value of K if
2 NO2(g) N2O4(g) at 100o C
What is K
if coefficients are doubled?
 If we multiply the equation by a
constant
2jA + 2kB
2lC + 2mD
 Then the equilibrium constant is
K’ =
[C]2l[D]2m = ([C]l[D]m)2 = K2
[A]2j[B]2k ([A] j[B]k)2
What is K
if coefficients are halved?
 If we multiply the equation by any
constant
njA + nkB
nlC + nmD
 Then the equilibrium constant is
K’ =
[C]nl[D]nm = ([C]l[D]m)n = Kn
[A]nj[B]nk
([A] j[B]k)n
5
What is the value of K if…
Given
½ I2(g) + ½ Cl2 (g)ICl(g)
Kc = 4.54 x 102 @ 25oC
What is the value of the equilibrium
constant for the reaction
2 I2(g) + 2 Cl2(g)  4 ICl(g)
The units for K
 Are determined by the various
powers and units of concentrations.
 They depend on the reaction.
 It is acceptable to write equilibrium
constants without units!!!
K is CONSTANT @ given T
 At any temperature.
Temperature affects rate.
 The equilibrium concentrations don’t
have to be the same, only K.
Equilibrium position is a set of
concentrations at equilibrium.
There are an unlimited number.
6
Equilibrium Constant
One for each Temperature
Calculate K
 N2 + 3H2
2NH3 at 25°C
At Equilibrium
 [N2]0 =1.000 M [N2] = 0.921M
 [H2]0 =1.000 M [H2] = 0.763M
 [NH3]0 =0 M
[NH3] = 0.157M
 Initial
Calculate K
 N2 + 3H2
2NH3 at 25°C
 Initial
At Equilibrium
 [N2]0 = 0 M
[N2] = 0.399 M
 [H2]0 = 0 M
[H2] = 1.197 M
 [NH3]0 = 1.000 M
[NH3] = 0.203M
 K is the same no matter what the
amount of starting materials
7
Equilibrium and Pressure for
rxns involving gases
 Some reactions are gaseous
 PV = nRT
 P = (n/V)RT
Equilibrium and Pressure
 2SO2(g) + O2(g)


Kp =
K=
2SO3(g)
(PSO3)2
(PSO2)2 (PO2)
[SO3]2
[SO2]2 [O2]
Homogeneous Equilibria
 So far every example dealt with
reactants and products where all
were in the same phase.
 We can use K in terms of either
concentration or pressure.
Units depend on reaction.
8
Heterogeneous Equilibria
 If the reaction involves pure solids or
pure liquids the concentration of the
solid or the liquid doesn’t change.
 As long as they are not used up we
can leave them out of the equilibrium
expression.
Write the Equilbrium
Expression
 H2(g) + I2(s)
K=
2HI(g)
[HI]2
[H2][I2]
 But the concentration of I2 does not
change; delete it!
 K = [HI]2
[H2]
Write the equilibrium constant for the
heterogeneous reaction
9
The Reaction Quotient, Q
 Tells you the direction the reaction
needs to go to reach equilibrium
 Calculated the same as the
equilibrium constant, but for a system
not at equilibrium
 Q = [Products]coefficient
[Reactants] coefficient
 Compare value to equilibrium constant
What Q tells us
 If Q<K
Not enough products
Shift to right
 If Q>K
Too many products
Shift to left
 If Q=K system is at equilibrium
Example
for the reaction
2NOCl(g)
K = 1.55 x
10-5 M
2NO(g) + Cl2(g)
at 35ºC
In an experiment 0.10 mol NOCl, 0.0010 mol
NO(g) and 0.00010 mol Cl2 are mixed in 2.0 L
flask. Which direction will the reaction proceed
to reach equilibrium?
10
Le Châtelier’s Principle
 If a stress is applied to a system at
equilibrium, the position of the
equilibrium will shift to reduce the
stress.
 3 Types of stress
Concentration
Pressure
Temperature
LeChatelier
Change amounts of reactants/ products
 Adding product makes Q>K
 Removing product makes Q<K
 Removing reactant makes Q>K
 Adding reactant makes Q<K
Determine the effect on Q, will tell you
the direction of shift
LeChatelier: Change Pressure
 Change Pressure by changing volume
 Decrease Volume…Increase Pressure
System will move in the direction that has the
least moles of gas; “SHIFT”
Because partial pressures (and
concentrations) change, a new equilibrium
must be reached.
System tries to minimize the moles of gas if
volume is reduced
 Decrease Pressure?
11
LeChatelier: Change in Pressure
 By adding an inert gas
 Partial pressures of reactants and
product are not changed
 No effect on equilibrium position
 No “SHIFT”
LeChatelier: Change in Temperature
 Affects the rates of both the forward
and reverse reactions.
 Doesn’t just change the equilibrium
position, changes the equilibrium
constant, new K Value
 The direction of the shift depends on
whether it is exo- or endothermic
Exothermic
 H<0
Releases heat
Think of heat as a product
Case 1: Raising temperature
Increase in product
Reaction shifts to “get rid of excess”
push toward reactants
Shifts to left
Case 2: Decrease temperature ?
12
Endothermic
 H>0
Absorbs heat
Think of heat as a reactant
Case 1: Raising temperature
Increase in reactant
Reaction shifts to “get rid of excess”
push toward product
Shifts to right
Case 2: Decrease temperature ?
Solving Equilibrium Problems
What if you’re not given
equilibrium concentration?
 Set up ICE Table
 Look at initial concentrations, compare Q
and K to determine which direction
reaction will shift (change row)
 Set up equilibrium expression, solve for x
 Solve for all equilibrium concentrations
13
Equilibrium Problem Example
H2(g) + I2(g)
2HI(g)
2
K = 7.1 x 10 at 25ºC
Calculate the equilibrium concentrations if a
5.00 L container initially contains 15.8 g of H2
294 g I2 .
[H2]0 = (15.8g/2.02)/5.00 L = 1.56 M
[I2]0 = (294g/253.8)/5.00L = 0.232 M
[HI]0 = 0
SET UP ICE TABLE.
COMPARE Q TO K FOR CHANGE ROW
 Q= 0;
Q<K so more product will be formed.
Set up table of initial, change in and
equilibrium concentrations.
 Assumption since K is large- reaction will
almost go to completion.
 Stoichiometry tells us I2 is LR, it will be
smallest at equilibrium let it be x
H2(g) + I2(g)
initial
1.56 M 0.232 M
change
-x
-x
equilibrium 1.56 -x .232 – x
2 HI(g)
0M
+x
x
Set up Equilibrium Expression, SOLVE
FOR X
14
H2(g) + I2(g)
initial
1.56 M 0.232 M
change
-x
-x
equilibrium 1.56 -x .232 – x
2 HI(g)
0M
+x
x
Plug values into equilibrium expression
H2(g) + I2(g)
initial
1.56 M 0.232 M
change
-x
-x
equilibrium 1.56 -x .232 – x
2 HI(g)
0M
+x
x
Use solver to solve for value of x
Plug X value into each to find
concentrations of each substance at
equilibrium
X = 2.3 x 10-4
find the other concentrations
 I2 =
2.3 x 10-4 M
 H2 = 1.328 M
 HI = 0.464 M
15
Equilibrium Problem Example
Cl2(g) + O2(g)
2ClO(g)
K = 156
In an experiment 0.100 mol ClO, 1.00 mol O2 and
0.0100 mol Cl2 are mixed in a 4.00 L flask.
If the reaction is not at equilibrium, which way will
it shift?
Calculate the equilibrium concentrations.
Cl2(g) + O2(g)  2ClO(g)
K = 156
I
C
E
At an elevated temperature, the reaction:
has a value of Keq = 944.
If 0.234 mol IBr is placed in a 1.00 L. flask and
allowed to reach equilibrium, what is the
equilibrium concentration in M. of I2?
16
I
C
E
Equilibrium Problem Example
2NOCl
2NO +Cl2
K= 1.6 x 10-5
If 1.20 mol NOCl, 0.45 mol of NO and 0.87 mol
Cl2 are mixed in a 1.00 L container
What are the equilibrium concentrations?
I
2NOCl
2NO +
1.20
0.45
Cl2
0.87
C
E
X= 8.2 x 10-3
17
Equilibrium Problem Example
For the reaction Cl2 (g) + O2 (g)
2ClO(g)
K = 6.4 x 10-3
In an experiment 0.100 mol ClO(g), 1.00 mol O2 and
1.00 x 10-2 mol Cl2 are mixed in a 4.00 L container.
What are the equilibrium concentrations?
Cl2 (g) + O2 (g)
I
C
E
2ClO(g)
Equilibrium Problem Example
H2(g) + I2 (g)
2HI(g)
K=38.6
What are the equilibrium concentrations if 1.800
mol H2, 1.600 mol I2 and 2.600 mol HI are mixed
in a 2.000 L container?
18
Problems Involving Pressure
For the reaction N2O4(g)
2NO2(g)
KP = .131 atm. What are the equilibrium
pressures if a flask initially contains 1.000 atm
N2O4?
19