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Le Chatelier's Principle

Key Concepts

In 1884, the French Chemist Henri Le Chatelier suggested that equilibrium systems tend to compensate for the effects of perturbing influences.

When a system at equilibrium is perturbed (or disturbed), the equilibrium position will shift in the direction which tends to minimise, or counteract, the effect of the disturbance.

For the system: reactants products     ΔH = ? kJ mol-1

If the equilibrium is perturbed (disturbed) by an increase in

  • reactant concentration in a solution:
    equilibrium position shifts right, more products produced

  • product concentration in a solution:
    equilibrium position shifts left, more reactants produced

  • volume available to a gas mixture (decrease in pressure):
    equilibrium position shifts to side with most gas molecules

  • temperature when ΔH is positive (endothermic)
    equilibrium position shifts right, more products produced

  • temperature when ΔH is negative (exothermic)
    equilibrium position shifts left, more reactants produced

Changes in Concentration of Aqueous Solutions

Consider the following system at equilibrium at a constant temperature:
Fe3+(aq) + SCN-(aq)
(colourless)

 
FeSCN2+(aq)
(red)

All the species, that is all reactants and products, are in aqueous solution.

At equilibrium, the rate at which Fe3+(aq) and SCN-(aq) react to produce FeSCN2+(aq) is the same as the rate at which FeSCN2+(aq) breaks apart to produce Fe3+(aq) and SCN-(aq).

The equilibrium position will be determined by

  • the concentration of each reactant, [Fe3+(aq)] and [SCN-(aq)], and the product, [FeSCN2+(aq)]

  • the temperature of the system

The equilibrium position will NOT be effected by

  • changes in volume or pressure because no gas species are present.

  • addition of a catalyst (a catalyst speeds up the rate of the forward and reverse reactions equally and does not change the equilibrium position)

Now consider what happens when the equilibrium position is perturbed (or disturbed) by changing the concentration of reactants or products while maintaining a constant temperature:

(a) Increasing the Concentration of a Reactant (at Constant Temperature)


Fe3+(aq) + SCN-(aq)
(colourless)

 
FeSCN2+(aq)
(red)
Perturbation: Addition of some other soluble salt of Fe3+
Immediate impact: Increase in concentration of Fe3+(aq)
Application of Le Chatelier's principle: Equilibrium position moves to the right, using up the some of the additional reactants and produces more FeSCN2+(aq).
New Equilibrium Position Established: Solution becomes a darker red colour because of the increase in concentration of FeSCN2+(aq).

(b) Decreasing the Concentration of a Reactant (at Constant Temperature)


Fe3+(aq) + SCN-(aq)
(colourless)

 
FeSCN2+(aq)
(red)
Perturbation: Removal of some of the Fe3+(aq), for example by precipitation.
Immediate impact: Decrease in the concentration of the reactant Fe3+(aq)
Application of Le Chatelier's principle: Equilibrium position moves to the left to produce more Fe3+(aq)
New Equilibrium Position Established: Solution becomes less red as concentration of FeSCN2+(aq) decreases as it is consumed to make reactants.

(c) Increasing the Concentration of a Product (at Constant Temperature)


Fe3+(aq) + SCN-(aq)
(colourless)

 
FeSCN2+(aq)
(red)
Perturbation: Addition of some other soluble salt of FeSCN2+
Immediate impact: Solution becoming a darker red as the concentration of FeSCN2+(aq) increases
Application of Le Chatelier's principle: Equilibrium position moves to the left to use up some of the additional FeSCN2+(aq) and produces more Fe3+(aq) and SCN-(aq)
New Equilibrium Position Established: Solution becomes less red than it was immediately following the addition of the FeSCN2+ as the concentration of FeSCN2+(aq) decreases as it is consumed in order to re-make reactants.

(d) Decreasing the Concentration of a Product (at Constant Temperature)


Fe3+(aq) + SCN-(aq)
(colourless)

 
FeSCN2+(aq)
(red)
Perturbation: Removing some of the FeSCN2+(aq), for example by precipitation
Immediate impact: Solution losing some of its colour due to decreased concentration of FeSCN2+(aq)
Application of Le Chatelier's principle: Equilibrium position will move to the right to produce more FeSCN2+(aq)
New Equilibrium Position Established: Solution becomes a darker red colour than it was immediately following the removal of the FeSCN2+(aq) as the concentration of FeSCN2+(aq) increases as it is produced.

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Gaseous Solutions and Changes in Concentration

Consider the following gaseous system at equilibrium at constant temperature and constant volume:

2NO2(g)
(red-brown)

 
N2O4(g)
(colourless)

All the species, reactant and product species, are gases.

At equilibrium, the rate at which NO2 molecules break apart to form N2O4 molecules is the same as the rate at which N2O4 break apart to make NO2 molecules.

The equilibrium position will be determined by the

  • concentration of the gases (the amount of each gas in a given volume)

  • temperature

  • volume of the vessel since concentration is the amount of gas per unit volume, changing the volume occupied by a gas will change its concentration
    (Note that changing the volume of the vessel holding the gases will also change the pressure inside the vessel since volume is inversely proportional to pressure by Boyle's Law)

The equilibrium position will NOT be effected by the addition of

  • a catalyst because a catalyst will speed up the rate of the forward and reverse reactions equally

  • an inert gas while maintaining constant volume because this does not effect the concentration of gaseous reactants and gaseous products

Now consider what happens when the equilibrium position is perturbed (disturbed) by changing the concentration of reactants or products at constant temperature and volume:

(a) Increasing the Concentration of a Gaseous Reactant (at Constant Temperature and Volume)


2NO2(g)
(red-brown)

 
N2O4(g)
(colourless)
Perturbation: Addition of more molecules of NO2(g) while maintaining the system at constant temperature and volume.
Immediate impact: Increase in concentration of NO2(g) therefore increase in red-brown colour.
(Increase in partial pressure of NO2(g))
Application of Le Chatelier's principle: Equilibrium position moves to the right, using up some of the additional NO2(g) and produces more N2O4(g).
New Equilibrium Position Established: Solution becomes a lighter red-brown colour than it was immediately following the addtion of NO2(g) as the concentration of NO2(g) decreases as it is consumed.

(b) Decreasing the Concentration of a Gaseous Reactant (at Constant Temperature and Volume)


2NO2(g)
(red-brown)

 
N2O4(g)
(colourless)
Perturbation: Removal of some molecules of NO2(g) while maintaining the system at constant temperature and volume.
Immediate impact: Decrease in the concentration of the reactant NO2(g) therefore reduction in the red-brown colour of the mixture.
(Decrease in partial pressure of NO2(g))
Application of Le Chatelier's principle: Equilibrium position moves to the left to produce more NO2(g)
New Equilibrium Position Established: Solution becomes more red-brown than it was immediately following the removal of NO2(g) as the concentration of NO2(g) increases as it is produced

(c) Increasing the Concentration of a Product (at Constant Temperature and Volume)


2NO2(g)
(red-brown)

 
N2O4(g)
(colourless)
Perturbation: Addition of more molecules of N2O4 while maintaining the system at constant temperature and volume.
Immediate impact: Concentration of N2O4 increases.
(Increase in partial pressure of N2O4)
Application of Le Chatelier's principle: Equilibrium position moves to the left to use up some of the additional N2O4(g) and produces more NO2(g)
New Equilibrium Position Established: Solution becomes a deeper red-brown colour as the concentration of NO2(g) increases as it is produced.

(d) Decreasing the Concentration of a Product (at Constant Temperature and Volume)


2NO2(g)
(red-brown)

 
N2O4(g)
(colourless)
Perturbation: Removing some molecules of N2O4 while maintaining the system at constant temperature and volume.
Immediate impact: Decrease in concentration of N2O4(g)
(Decrease in partial pressure of N2O4)
Application of Le Chatelier's principle: Equilibrium position will move to the right to produce more N2O4(g)
New Equilibrium Position Established: Solution becomes a lighter red-brown colour as the concentration of NO2(g) decreases as it is consumed.

(e) Addition of an Inert Gas (at Constant Temperature and Volume)


2NO2(g)
(red-brown)

 
N2O4(g)
(colourless)
Perturbation: Addition of an inert gas such as argon while maintaining the system at constant temperature and volume.
Immediate impact: No change in concentration of NO2(g) and no change in concentration of N2O4(g).
(No change in partial pressure of NO2(g) or N2O4)
Application of Le Chatelier's principle: Equilibrium position does not change.
New Equilibrium Position Established: Solution remains the same colour because there is no change in the concentration of NO2(g)

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Changes in Volume of a Gaseous System at Equilibrium

Consider the following gaseous system at equilibrium:

2NO2(g)
(red-brown)

 
N2O4(g)
(colourless)

All the species, reactant and product species, are gases.

At equilibrium, the rate at which NO2 molecules break apart to form N2O4 molecules is the same as the rate at which N2O4 break apart to make NO2 molecules.

The equilibrium position will be determined by the

  • concentration of the gases (the amount of each gas in a given volume)

  • temperature

  • volume of the vessel since concentration is the amount of gas per unit volume, changing the volume occupied by a gas will change its concentration
    (Note that changing the volume of the vessel holding the gases will also change the pressure inside the vessel since volume is inversely proportional to pressure by Boyle's Law)

The equilibrium position will NOT be effected by the addition of

  • a catalyst because a catalyst will speed up the rate of the forward and reverse reactions equally

  • an inert gas while maintaining constant volume because this does not effect the concentration of gaseous reactants and gaseous products (see section above)

Now consider what happens when the equilibrium position is perturbed (disturbed) by changing the volume of the vessel containing the gaseous reaction mixture while maintaining a constant temperature:

(a) Reducing the Volume of the Reaction Vessel (at constant temperature)


2NO2(g)
(red-brown)

 
N2O4(g)
(colourless)
Perturbation: Reducing the volume of the reaction vessel, for example by depressing the plunger of a syringe, while maintaining a constant temperature.
Immediate impact: Concentration of both NO2(g) and N2O4 increases
Total gas pressure inside reaction vessel increases.
Application of Le Chatelier's principle: Equilibrium position shifts to the right, the side with the fewest gas molecules, to reduce the total number of gas molecules in the vessel and thereby decrease the pressure inside the vessel.
New Equilibrium Position Established: Solution becomes a lighter red-brown colour because the concentration of NO2(g) decreases as it is consumed.

(b) Increasing the Volume of the Reaction Vessel (at constant temperature)


2NO2(g)
(red-brown)

 
N2O4(g)
(colourless)
Perturbation: Increasing the volume of the reaction vessel, for example by pulling the plunger of a syringe up, while maintaining a constant temperature.
Immediate impact: Concentration of both NO2(g) and N2O4 decreases
Total gas pressure inside reaction vessel decreases.
Application of Le Chatelier's principle: Equilibrium position shifts to the left, the side with the most gas molecules, to increase the total number of gas molecules in the vessel and thereby increase the gas pressure inside the vessel.
New Equilibrium Position Established: Solution becomes a deeper red-brown colour because the concentration of NO2(g) increases as it is produced.

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Changes in Temperature

In an endothermic reaction, energy can be considered as a reactant of the reaction.

In an exothermic reaction, energy can be considered as a product of the reaction.

Le Chatelier's Principle predicts that the equilibrium position will shift in order to:

  • consume more heat if the reaction mixture is heated
    (that is, the endothermic reaction is favoured when the reaction mixture is heated)

  • produce more heat if the reaction mixture is cooled
    (that is, the exothermic reaction is favoured when the reaction mixture is cooled)

Endothermic Reactions at Equilibrium

Consider the following reaction at equilibrium:
H2(g)
(colourless)
+
 
I2(s)
(purple)

 
2HI(g)
(colourless)
    ΔH = +52 kJ mol-1

This reaction can also be written with the energy term incorporated into the equation on the side with the reactants:

H2(g)
(colourless)
+
 
I2(s)
(purple)
+ 52 kJ
 

 
2HI(g)
(colourless)

When the system is at equilibrium, the rate at which H2(g) combines with I2(s) to produce HI(g) is the same as the rate at which HI(g) breaks apart to form H2(aq) and I2(s).

The equilibrium position will be determined by the

  • concentration of H2(g) and HI(g)

  • temperature

  • volume of the reaction vessel (or total pressure since volume is inversely proportional to pressure by Boyle's Law) since gaseous species are present

The equilibrium position will NOT be effected by the

  • addition of a catalyst (a catalyst speeds up the rate of the forward and reverse reactions equally and does not change the equilibrium position)

  • addition of an inert gas while maintaining constant volume because this has no effect on the concentration of gaseous reactants and products.

  • amount of I2(s) because the equilibrium position is not dependent on amounts of species present but rather on the concentration of species present.

Now consider what happens when the equilibrium position is perturbed (disturbed) by changing the temperature of the system:

(a) Increasing the Temperature of the System


H2(g)
(colourless)
+
 
I2(s)
(purple)
+ 52 kJ
 

 
2HI(g)
(colourless)
Perturbation: Heating the reaction mixture.
Immediate impact: Temperature of the system increases.
Application of Le Chatelier's principle: Equilibrium position shifts to the right in order to consume some of this additional heat energy to compensate for the heat gained.
New Equilibrium Position Established: More H2(g) and I2(s) will be consumed so there will be less purple solid (I2(s)) present in the vessel.
Temperature of the system does not increase as much as it would have otherwise.

(b) Decreasing the Temperature of the System


H2(g)
(colourless)
+
 
I2(s)
(purple)
+ 52 kJ
 

 
2HI(g)
(colourless)
Perturbation: Removing heat by cooling the reaction mixture.
Immediate impact: Temperature of the system decreases.
Application of Le Chatelier's principle: Equilibrium position shifts to the left in order to produce additional heat energy to compensate for the lost heat.
New Equilibrium Position Established: More H2(g) and I2(s) will be produced in the reaction vessel so there will an increase in the amount of purple solid (I2(s)) in the vessel.
Temperature of the system does not decrease as much as it would have otherwise.

Exothermic Reactions at Equilibrium

Consider the following system at equilibrium:

Ag+(aq) + Cl-(aq)
(colourless)

 
AgCl(s)
(white)
    ΔH = -112 kJ mol-1
 

This reaction can also be written with the energy term incorporated into the equation on the side with the the products:

Ag+(aq) + Cl-(aq)
(colourless)

 
AgCl(s)
(white)
+ 112 kJ
 

When the system is at equilibrium, the rate at which Ag+(aq) combines with Cl-(aq) to produce a precipitate of AgCl(s) is the same as the rate at which AgCl(s) breaks apart to form Ag+(aq) and Cl-(aq).

The equilibrium position will be determined by the

  • concentration of Ag+(aq) and Cl-(aq)

  • temperature

The equilibrium position will NOT be effected by the

  • addition of a catalyst (a catalyst speeds up the rate of the forward and reverse reactions equally and does not change the equilibrium position)

  • addition of an inert gas while maintaining constant volume, or change in volume of vessel because no gas species are present

  • amount of AgCl(s) because the equilibrium position is not dependent on amounts of species present but rather on the concentration of species present

Now consider what happens when the equilibrium position is perturbed (disturbed) by changing the temperature of the system:

(a) Increasing the Temperature of the system.


Ag+(aq) + Cl-(aq)
(colourless)

 
AgCl(s)
(white)
+ 112 kJ
 
Perturbation: Heating the reaction mixture.
Immediate impact: Temperature of the system increases.
Application of Le Chatelier's principle: Equilibrium position shifts to the left in order to consume some of the additional heat energy.
New Equilibrium Position Established: White AgCl(s) will be consumed so there will be less AgCl(s) in the vessel.
Concentration of Ag+(aq) and of Cl-(aq) increases.
Temperature of the system does not increase as much as it would have otherwise.

(b) Decreasing the Temperature of the system.


Ag+(aq) + Cl-(aq)
(colourless)

 
AgCl(s)
(white)
+ 112 kJ
 
Perturbation: Removing heat by cooling the reaction vessel.
Immediate impact: Temperature of the system decreases.
Application of Le Chatelier's principle: Equilibrium position shifts to the right in order to produce more heat energy.
New Equilibrium Position Established: White AgCl(s) will be produced so there will be more AgCl(s) in the vessel.
Concentration of Ag+(aq) and of Cl-(aq) decreases.
Temperature of the system does not decrease as much as it would have otherwise.

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see also Solubility and Le Chatelier's Principle
see also Henry's Law (gas solubility)


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