Slide 1
Chang Chapter 14
Read Sections 14.1 to 14.5
Chemical Equilibrium
An Outline
• Understand the dynamic equilibrium between
reactants and products in a chemical reaction
• Use of equilibrium constants to predict the concentrations of reactants and products
• Apply LeChatelier’s Principle
– How the knowledge of chemical equilibrium allows us to use temperature and pressure to push a
reaction towards a desired outcome, e.g., maximize yield in the synthesis of NH 3, starting from the
reactants N2 and H2
Dynamic Equilibrium
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Equilibrium – two opposing processes taking place at equal rates.
Physical Change
NaCl(s)
NaCl (aq)
Chemical Change (or reaction)
N2(g) + 3 H2(g)  2 NH3(g)
Ammonia: An important compound
• Roughly 1010 kg of ammonia produced every year in USA alone
• Worldwide, 80% of the amount used in fertilizers
• In the early 1900’s there was a pressing need for synthesizing NH3
The Problem when making NH3
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It looks like there is an abundant supply of N2 and H2
To do the reaction (an example of nitrogen fixation)
N2(g) + 3 H2(g)  2 NH3(g)
However,the actual yield is very limited at 25 C and 1 atm
Why? One has to break the triple bond first in N2
To “copy” nitrogen fixing bacteria on an industrial scale is complicated
To increase the actual yield, Haber applied principles of chemical equilibria, which we
will examine in Chapter 14
Haber-Bosch process
• Haber and Bosch devised a synthesis of ammonia from N2 and H2 that
takes place
• at low temperature
• high pressure
• uses a catalyst
• Nobel prize in Chemistry
Manufacture of NH3
Equilibrium: An Example
• Suppose we place N2 and H2 gases in a reactor
• The concentrations of N2 and H2 decrease with time, and eventually attain a constant
value
• The concentration of NH3 is initially zero, it increases with time and eventually
achieves a constant value which under normal conditions is small.
The meaning of “equilibrium”
• When the concentrations of reactants and products have stopped changing, we say that
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the system has achieved an equilibrium
– Note: The reactants are still there!
The reactants react to form the product or the product falls apart.
The forward and reverse reactions occur at equal rates
The amounts of each of the species remains constant at equilibrium, if we don’t disturb
the system
Notation: N2(g) + 3 H2(g)
2 NH3(g)
double arrows = reversible system
Forward and reverse reaction
Equilibrium is Dynamic
Picture slide
Equilibrium Constant
• Consider the reaction
• 2 SO2(g) + O2(g)
2 SO3 (g)
• Suppose we take different amounts of each of the three gases and
allow them to react at 1000 K. Let’s look at numbers.
• (This reaction is important in the synthesis of sulfuric acid)
Equilibrium constant
The law of mass action
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(We will first consider equilibria involving gases only, then generalize)
For a reaction involving gases A, B, C, D
aA+bB
cC +dD
where a, b, c, d are stoichiometric coefficients,
We can calculate the equilibrium constant
Kc = [C]c[D]d
[A]a[B]b
Things to know about Kc
Equilibrium constants
• [C] denotes equilibrium concentration of C
1 mol/L
• We do this so that the equilibrium constant turns out to be a unitless number
In Kc, c stands for concentration. We can also use partial pressures to calculate
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equilibrium constants, denoted Kp
• To get Kc, divide product concentrations by reactant concentrations, with each
concentration raised to the power of appropriate stoichiometric coefficient
• Kc =
Practice Exercises
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Write the expression for the equilibrium constant Kc for each
of the following reactions
1. N2(g) + 3 H2(g)
2 NH3(g)
2. CO(g) + Cl2(g)
COCl2(g)
3. N2O4 (g)
2 NO2(g)
Interpreting equilibrium constants
Three possible cases:
K<<1
K=1
K >>1
Pictures missing
Heterogeneous Equilibria
• Homogeneous equilibria
• All reactants and products are in the same phase, for example
– Gases as in the synthesis of ammonia
– Species dissolved in water, e.g., acetic acid dissociating in water
• Heterogeneous Equilibria
• Reactants and products are in more than one phase or state
• CaCO3 (s)
CaO (s) + CO2 (g)
Examples of Heterogeneous Equilibria
• This reaction used both solids and gases
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Ni(s) + 4 CO(g)
Ni(CO)4(g)
• The equilibrium is not affected so long as there is some solid Ni present
• We ignore (leave out) Ni(solid) in the expression for equilibrium constant
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Kc = [Ni(CO)4]
[CO]4
Practice Exercises
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Write the expression for the equilibrium constant (Kc) for each
of the following:
1. NH4S(s)
NH3(g) + H2S(g)
2. 2 KClO3(s)
2 KCl(s) + 3 O2(g)
3. CaCO3(s)
CaO(s) + CO2(g)
Heterogeneous Equilibria
• Recall that the concentration of a solid (or liquid) is not used while writing the
expression for Kc in a system that has gases as reactants and/or products
• In a heterogeneous system involving gases and a solid or liquid, addition of solid or
liquid reactants or products does not affect the equilibrium
• Ni(s) + 4 CO(g)
Ni(CO)4(g)
• So long as there is some Ni(s) present, adding or removing Ni(s) does not affect the
equilibrium
Two ways of specifying amounts of a gas
• For a system involving gases, we can specify the amount of a gas as either
– concentration (moles/L)
– partial pressure (atm)
• Recall: the partial pressure of a gas in a mixture is the pressure that each gas
would exert if it occupied the container alone
• Of course, concentration and partial pressure of a gas can be related using the
ideal gas law
Introducing Kp
• For a reaction involving gases A, B, C, D
• aA+bB
cC +dD
• where a, b, c, d are stoichiometric coefficients,
We can calculate the equilibrium constant
Kp = pcc pDd
pAa pBb
• where pA is unitless and has the magnitude of the partial pressure in atmospheres of
the species A
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The form is still pressures of products over reactants, raised to powers of
stoichiometric coefficients
Practice Exercises
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Write the expression for the equilibrium constant (Kp) for each
of the following:
1. N2(g) + 3 H2(g)
2 NH3(g)
2. NH4S(s)
NH3(g) + H2S(g)
3. 2 KClO3(s)
2 KCl(s) + 3 O2(g)
Direction of Reaction
• Suppose we are carrying out a reaction, and we wish to know what the
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direction of the reaction is
There are three possibilities
The reaction is at equilibrium
There is a tendency to produce more products (go forward)
There is a tendency to make more reactants
(go backwards)
Three Approaches to Equilibrium
• Expt 1: 1.00 M CO & H2 ; zero CH3OH
Pictures missing
Direction of Reaction
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Calculate Q, the reaction quotient
Q has the same form as K except that we use it at any stage of the reaction
If Q = K, equilibrium
If Q < K, forward
If Q> K, backward
Practice Exercises
• For the reaction
2 SO2(g) + O2 (g)
Kc = 1.7 x 106 K
2 SO3 (g)
• Suppose we take 1.20 x 10-3 mol SO2(g), 5.0 x 10-4 mol O2(g)
and 1.0 x 10-4 mol SO3 (g) in a 500.0 mL container.
• Calculate Qc.
• Will more SO3(g) tend to form?
Disturbing a system at equilibrium
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Suppose we take a system that is at equilibrium, and disturb it by
Adding or removing reactants or products
Change in Pressure or changing the volume of the container
Changing the temperature
Le Chatelier’s principle: When a stress is applied to a system in dynamic
equilibrium, the equilibrium tends to adjust to minimize the effect of the stress
Addition of Reactants:
Homogeneous System
Picture missing
Addition of Products:
Homogeneous System
Picture missing
An Illustration
• Consider the reaction N2O4(g)
2 NO2(g)
Kc = 0.20 at 100° C.
• Suppose we have a mixture of NO2 and N2O4 gases in equilibrium
• If we add N2O4, what should happen?
• more NO2 will form
• Conversely, if we add NO2 to a system initially at equilibrium, N2O4 will
form
How does this work?
• Recall Qc (the reaction quotient)
• It has the form (product concentrations) divided by (reactant
concentrations), each raised to an appropriate power
• At equilibrium, Qc = Kc
• If a reactant is added, Qc falls below the value of Kc, and the reaction goes
forward until Qc equals Kc and an equilibrium is achieved again
Playing with actual numbers
• Consider the reaction N2O4(g)
2 NO2(g)
Kc = 0.20 at 100° C.
• Suppose we have a 1.00 L container, 0.447 mole of NO2 and 1.0 mole of
N2O4
• Are we at equilibrium?
• Now add 1.0 mole of N2O4
• Kc is still unchanged, it depends only on temperature
• Now calculate Qc
• In what direction will the reaction proceed?
LeChatelier’s Principle
Effect of Change in concentration
Data and picture missing
Determine what happens if you add 1.0 mole SO3 to
the 10.0 L bulb
To Calculate the K eq
Effect of Change of Volume
Effect of decreasing Volume or increasing
pressure
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Consider the reaction
 N2(g) + 3 H2(g)
2 NH3(g)
Suppose the system is initially at equilibrium
The system is compressed by increasing pressure
The larger the number of gaseous molecules, larger the pressure (PV=
nRT)
Reduce the number of moles of gaseous species
Shifts to the RIGHT to FORM MORE NH3 (Go from 4 reactant moles to
2 product moles)
Effect of Change in Pressure or Change in
Volume
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Decrease in volume or increase in pressure of a reaction mixture of gases
at equilibrium drives the reaction in the direction that reduces the number
of gas phase molecules
There are two instances when increasing the pressure does not affect the
system at equilibrium
The number of moles of gaseous reactants equals the number of gaseous
products as in the reaction,
 2 HI(g)
H2(g) + I2(g)
An inert gas is added (no shift in equilibrium since the partial pressures
are unaffected)
Le Chatelier's Principle
Effect of Temperature
• For an endothermic reaction the rate constant increases with increasing temperature
• For an endothermic reaction at equilibrium, an increase in temperature favors the
formation of products
• Why? Think in terms of Q
• Can you predict the effect of increasing the temperature for an exothermic reaction at
equilibrium?
Effect of Temperature on Equilibrum
Le Chatelier's Principle
Consider the equilibrium
AgCl (s)
Ag +(aq) + Cl _(aq)
What would be the effect of
a. adding more silver chloride solid?
Hint: Can more AgCl solid dissolve if nothing else was changed?
b. adding more water
c. increasing the temperature
Effect of a Catalyst on Equilibrium
Effect of Catalyst
Graph of Ea missing
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How does lowering Ea affect the rate of a reaction?
How does lowering Ea affect the position of equilibrium?
Simultaneous Equilibria
The magnitude of K
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Caution! The magnitude of K, the equilibrium constant, can be
changed only by changing temperature
• Its numerical value cannot be changed by
A. Adding reactants/products
B. Compressing/Expanding the volume of the reactor
C. Adding a catalyst