Chemistry 12
Chemical Equilibrium: Factors Affecting Equilbrium
Le Chatelier’s Principle
To an equilibrium system, a stress (are you stressed?) is any change that
has a different effect on either the forward or reverse reaction rates, thus
disrupting the equilibrium. Provided that the “new” conditions remain constant, a
new state of equilibrium will be reached.
The French chemist Henri Le Chatelier described the behavior of equilibrium
systems with a general principle known as (whatelse?) Le Chatalier’s Principle.
“When a stress is applied to a system already at equilibrium, the system
readjusts so as to relieve or offset the stress.”
A simple paraphrase: “Whatever we do, nature tends to undo”.
So the effects of varying temperature, concentrations and pressure conditions of
a system at equilibrium can be predicted using Le Chatelier’s Principle.
Consider the following equilibrium system:
2NO(g) + Cl2(g)

2NOCl(g) + 76 kJ
1. Temperature Change
If the temperature of the equilibrium system is decreased and maintained
at this new lower temperature, Le Chatelier’s Principle tells us that when
the amount of heat energy is decreased, the reaction “shifts’ so as to
produce more heat. Since this reaction is exothermic, the reaction shifts to
the product (right) side so as to produce more heat energy. Also, more
NOCl is produced.
In terms of chemical kinetics: as temperature is decreased, both the
forward and reverse reaction rates decrease, but the reverse
(endothermic direction) reaction rate decreases more that the rate of the
forward reaction. A net forward (exothermic) reaction results. This is
always true, because the endothermic reaction generally has a harder task
(due to a higher Ea) and so will suffer more from being hindered
(decreased temperature).
NOTE: Keq is temperature dependent – other conditions such as
concentration and volume changes do not affect the value of Keq.
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A concentration-time graph illustrates the effect temperature has on an
equilibrium system:
2. Concentration change
If the [Cl2] increases, Le Chatelier’s Principle predicts that the reaction will
shift so as to partially decrease the amount of added Cl2. As a result, the
equilibrium shifts to the right, producing more NOCl.
In terms of chemical kinetics: An increase in [Cl2] increases forward rate
but has no immediate effect on the reverse rate. This results in a net
forward reaction (shifts right). The system then re-equilibrates in the
same manner that it established equilibrium in the first place – i.e. as the
net forward reaction proceeds, [reactant] falls and thus so thus forward
reaction rate decreases too. Simultaneously, [product] rises and hence
reverse rate increases. This continues until the two opposing rates are
equal again and equilibrium is re-established.
Graphically:
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3. Volume change
Since volume is related to pressure and molarity (concentration), a change
in volume can affect an equilibrium system. The concentrations of all
reactants and products can be changed simultaneously by changing the
volume of the reacting system. The volume of a gaseous system can be
changed by either compressing or expanding (decompressing) it. Note
that it is not possible for a shift to restore all of the chemicals’
concentrations, but some equilibria can shift to partially restore the total
chemical concentration.
Consider an increase in volume of a gaseous system (as above). Since
volume is inversely proportional to pressure (i.e. as volume increases,
pressure decreases and vice versa), the equilibrium will shift to the left
(reactant) side – the reason being that there are more moles of gas on
the left, and the pressure can only be increased by increasing the number
of moles of gas. Note also that the concentrations of all reactants and
products initially decrease as a result of the volume increase. (Why?)
So for 2NO(g) + Cl2(g)

2NOCl(g) the equilibrium will shift to the
left , producing more NO and Cl2. Since there are more moles of these
gases collectively (relative to NOCl), the pressure will then be partially
increased.
In terms of chemical kinetics: An increase in volume of a system means a
decrease in concentration of all species – as a result, the forward reaction
rate decreases and a net reverse reaction results (i.e. equilibrium shifts
left). Eventually, equilibrium is re-established.
Graphically:
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4. Addition of an inert gas
Adding an inert (noble) gas such as He at constant volume will increase
the total pressure in the system, but it has NO EFFECT on the
concentrations (or partial pressures) of reactants or products, since the
inert gas doesn’t participate in the reaction in any way.
5. Addition of a catalyst
A catalyst accelerates the reaction rate and helps the system reach
equilibrium faster. If the system is already at equilibrium, then both the
forward and reverse rates are equal. Addition of a catalyst would have no
effect on the equilibrium position since it will increase both the forward
and reverse reaction rates equally.
The Haber-Bosch Process: An Application of Le Chatelier’s Principle
Virtually all of the world’s ammonia (NH3) is produced via the Haber-Bosch
process and almost all of our inorganic nitrogen compounds (nitrates, nitrites
etc) are synthesized from this ammonia. About 10 million tons of nitric acid is
produced annually from ammonia. Nitric acid is used to produce ammonium
nitrate fertilizers, nitro explosives, plastics and dyes. In 1910, German chemist,
Fritz Haber developed the equipment and procedures for producing ammonia
from its corresponding elements (N2 and H2). He later received a Nobel Prize in
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chemistry for his work. Later, Carl Bosch further developed the process to the
industrial scale.
N2(g)
+
3H2(g)

2NH3(g)
H = -92.4 kJ
Consider a scenario where a chemist wishes to make NH3 in the lab by allowing a
single set of reactants achieve equilibrium in a closed container. An optimal
temperature and pressure would be ideal, so both percent yield and reaction rate
need to be considered.
(i)
To produce higher percent yields of NH3 at equilibrium, what
kind of temperatures would be needed (lower or higher)?
__________
(ii)
How would this affect the (forward and reverse) reaction
rates? ____________________________________
(iii)
To produce higher percent yields of NH3, what kind of pressure
would be needed (higher or lower)? ____________
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