Lesson 13.5 Le Chatelier's Principle
Suggested Reading

Zumdahl Chapter 13 Section 13.7
Essential Question

How is Le Chatelier's principle applied to chemical equilibria?
Learning Objectives:


State Le Chatelier's principle.
Apply Le Chatelier's principle to predict the qualitative effects of changes of temperature,
pressure, and concentration on the position of equilibrium and on the equilibrium
constant.
 State and explain the effect of a catalyst on an equilibrium reaction.
 Apply the concepts of kinetics and equilibrium to industrial processes.
Introduction
In 1884 the French chemist Henry-Louis Le Chatelier proposed one of the central concepts of
chemical equilibria. Le Chatelier's principle can be stated as follows: A change in one of the
variables that describe a system at equilibrium produces a shift in the position of the
equilibrium that counteracts the effect of this change.
Le Chatelier's principle describes what happens to a system when something momentarily takes
it away from equilibrium. This section focuses on three ways in which we can change the
conditions of a chemical reaction at equilibrium:
(1) changing the concentration of one of the components of the reaction
(2) changing the pressure on the system
(3) changing the temperature at which the reaction is run.
Changing Concentrations
To illustrate what happens when we change the concentration of one of the reactants or
products of a reaction at equilibrium, let's consider the following system at 500oC, N2(g) + 3H2(g)
⇌ 2NH3(g), Kc = 0.040. (Do you remember the name of the process described by this
equation?)
Concentration (M)
N2(g)
2NH3(g)
3H2(g) ⇌
Initial
0.100
0.100
0
Change
-x
- 3x
+2x
Equilibrium
0.099
0.097
0.0020
Note that the change in concentration is very small compared to the initial concentrations of
N2 and H2. This implies that very little ammonia is actually produced in this reaction. According
to this calculation, only 1% of the nitrogen is converted into ammonia.
What would happen if we increased the initial concentration of N2 by a factor of 10? The
reaction can't be at equilibrium any more because there is too much N2 in the system (Q < K).
Adding an excess of one of the reactants therefore places a stress on the system. The system
responds by minimizing the effect of this stress
by shifting the equilibrium toward the
products. The system will continue to make product until the equilibrium position is once again
equal to K.
Changing Pressure
The effect of changing the pressure on a gas-phase reaction depends on the stoichiometry of
the reaction. We can demonstrate this by looking at the result of compressing the following
reaction at equilibrium.
N2(g) + 3H2(g) ⇌ 2NH3(g)
Consider a system that initially contains 2.5 atm of N2 and 7.5 atm of H2 at 500oC, where Kp is
1.4 x 10-5 (What does the magnitude of K tell us about the reaction mixture?). If we allow the
reaction to come to equilibrium, and then compress the system by a factor of 10, we get the
following results.
Before Compression
After Compression
PNH3 = 0.12 atm
PN2 = 2.4 atm
PH2 = 7.3 atm
PNH3 = 8.4 atm
PN2 = 21 atm
PH2 = 62 atm
Before the system was compressed, the partial pressure of NH3 was only a small fraction of the
total pressure. After the system is compressed, the partial pressure of NH3 is almost 10% of the
total.
These data illustrate another application of Le Chatelier's principle. A reaction at equilibrium
was subjected to a stress
an increase in the total pressure on the system. The reaction then
shifted in the direction that minimized the effect of this stress. In this case, the reaction shifted
toward the products because this reduced the number of particles in the gas, thereby
decreasing the total pressure on the system.
Changing Temperature
Changing the concentrations of reactants or products shifts the position of the equilibrium, but
does not change the equilibrium constant for the reaction. Similarly, a change in the pressure on
a gas-phase reaction shifts the position of the equilibrium without changing the magnitude of the
equilibrium constant. Changes in temperature do change the magnitude of the equilibrium
constant for the reaction.
Chemical reactions either release heat to or absorb heat from their surroundings. If we consider
heat to be one of the reactants or products of a reaction, we can understand the effect of
changes in temperature on the equilibrium. Increasing the temperature of a reaction that gives
off heat is the same as adding more product. It places a stress on the reaction, which must be
alleviated by converting some of the products back to reactants.
Consider the reaction in which NO2 forms N2O4. This reaction is exothermic.
2NO2(g) ⇌ N2O4, ∆H = -57.20 kJ
Raising the temperature of this system is equivalent to adding excess product to the system,
except no product is actually added, so the equilibrium constant decreases as a result.
Similarly, if we add heat to and endothermic reaction, this is the same as adding additional
reactant, the equilibrium shifts to the right, and equilibrium constant increases. Since reactant is
not actually added, we have a change in the numerator of the equilibrium constant expression
without a corresponding change in the denominator, so K must change as a result.
Effect of a Catalyst
Recall that a catalyst is a substance that increases the rate of a reaction but is not consumed by
it. A catalyst has not effect on the position of equilibrium. However, it does affect the rate at
which the position of equilibrium is achieved. You learned in chapter 12 that a catalyst works by
lowering the activation energy. However, when a system moving towards equilibrium, the
catalyst lowers the activation energy for both the forward and the reverse reaction. The result is
that equilibrium is achieved more quickly.
Watch the following YouTube Video:
https://www.youtube.com/watch?v=lzq-31viF-0
Applying Concepts of Kinetics and Equilibrium to the Haber Process and the Contact
Process
This is an important learning objective within the IB curriculum. However, you should be able to
qualitatively apply the concepts of kinetics and equilibrium to any chemical reaction. For the AP
exam, make sure you can do so for the following reactions.
Haber Process for the production of ammonia, N2(g) + 3H2(g) ⇌ 2NH3(g).
The Contact Process for the production of sulfuric acid, 2SO2(g) + O2(g) ⇌ 2SO3(g).
Here are a couple of videos to help you learn this stuff.
Watch the following YouTube Video:
https://www.youtube.com/watch?v=3ifdYcyf54Y
Watch the following YouTube Video:
https://www.youtube.com/watch?v=3ifdYcyf54Y
Homework: Practice exercises 15.13, 15.14
Book questions page 617 questions 57-64