IB Chemistry home > Syllabus 2016 > Equilibrium > Changing conditions

Syllabus ref: 7.1

Le Chatelier's principle

This considers the effect of changing conditions on an equilibrium.

There is no fixed statement of the principle so it does not need to be learned as a definition.

One possible statement of the principle could be "Any system at equilibrium will respond to a change in conditions in such a way as to oppose that change."

Example: The fundamental concept is that if a system at equilibrium is able to reduce the effect of any applied change then it will do so.

You could consider the effect of pressure change on a rubber balloon. If the balloon is subject to increasing pressure, it will respond by expanding in an attempt to decrease the pressure once again.

Although this is not a chemical system, it demonstrates the idea very clearly.


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Changing conditions

The conditions that concern us are:

  1. Changing the amounts of reactants or products by addition or removal.
  2. Changing the pressure
  3. Changing the temperature

In each of these situations an equilibrium will respond, according to Le Chatelier's principle, in an attempt to reduce the effects of the change in conditions.

The process is as follows:

  1. 1 System at equilibrium
  2. 2 Change of conditions disturbs the equilibrium (i.e. system is now not at equilibrium)
  3. 3 System responds by moving forward or backwards as necessary to re-establish the equilibrium.

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Addition or removal of reactants

Addition of reactants (or products in the case of the reverse reaction) to a system at equilibrium disturbs that equiliibrium.

The system responds in such a way as to re-establish the equilibrium.

It does so by removing the added reactant, increasing the rate of the forward reaction and making more product until the equilibrium proportions are re-established.

This is an example of Le Chatelier's principle at work; the system counteracts the change in conditions, in this case the change in concentration of a reactant.

More mathematically, what is happening is that the ratio or products to reactants as given by the equilibrium law becomes not equal to the equilibrium constant. The system ceases to be at equilibrium and readjusts its reactant and product concentrations to re-establish the equilibrium.

Example: In the reversible reaction N2O4 2NO2, the system is at equilbrium when the concentration of N2O4 is 0.1 mol dm-3 and the concentration of NO2 is 0.4 mol dm-3. All of the reaction takes place in a 1 dm3 flask. What is the effect of adding 0.1 moles of N2O4 to the equilibrium?

The system products and reactants were originally in equilibrium and now more reactant has been added.

This change can only be opposed by making more product.

The system must now react to make more NO2 and less N2O4. It does so by moving the reaction in the forward direction until the equilibrium proportions are reestablished.

The general principle of Le Chatelier is followed and addition to the Left Hand Side moves the reaction to the Right Hand Side.

Example Add to the left push to the right. Take from the left pull to the left.


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Change of pressure

Changing the pressure of a system at equilibrium only has an effect on the reaction if there are different numbers of moles of gases on each side of the equilibrium.

Example: N2O4 2NO2 is affected by a change of pressure (1 mole on the left and 2 moles on the right hand side of the equilibrium)

2HI H2 + I2 is NOT affected by a change of pressure (2 moles on the left and 2 moles on the right hand side of the equilibrium)

The question remains, why?

In order to fully understand why change of pressure affects equilibria with unequal numbers of moles on either side, it is first important to understand how the pressure of a gas can be changed.

If the temperature is to be kept constant and the number of moles of gas is constant then the only way to change the pressure of a gas is to change the volume it occupies as PV = nRT (the ideal gas law). In the ideal gas law, if n, R and T are constants then only V can change P.

The definition of concentration is moles divided by volume. So if we are changing the volume, we are also changing the concentration. Now we can apply Le Chatelier's principle. If we increase the concentration on one side the reaction moves away from that side.

But, we are increasing the concentration on both sides! This is where the number of moles on each side is important.

If there are more moles on one side then an increase in concentration will have more effect on that side in the equilibrium law ratio.

Therefore, an increase in pressure moves the reaction away from the side with more moles and towards the side with fewer moles.

Le Chateliers principle once again shows the system responding to counteract a change in conditions.


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Change of temperature

Changing the temperature of a chemical reaction changes the rate of reaction. At equilibrium the rate of both the forward and back reactions are equal.

So how does a change in temperature affect the rate of forward and back reactions differently?

The answer is that it wouldn't if there were no energy change involved in the chemical reaction, i.e. if ΔH = 0 (not very likely).

If the reaction is exothermic then the activation energy for the forward reaction MUST be less than for the reverse reaction.

This means that an increase in temperature for the equilibrium has more of an effect on the rate of reaction involving the larger of the two activation energies, ie. the direction of endothermic change.

The mathematics are beyond the level of pre-university syllabi, but this can be demonstrated using the Arrhenius equation k = Ae-Ea/RT.

By Le Chatelier's principle the reaction responds to a change of conditions in such a way as to oppose that change.

If the temperature is increased, there is more energy available. The reaction absorbs this extra energy by moving in the direction of endothermic change.

Example: What is the effect of an increase in temperature on the equilibrium:

N2 + 3H2 2NH3   ΔH = -92 kJ mol-1

The forward reaction is exothermic ( ΔH is negative). The temperature is increased, there is more energy available therefore the reaction will move in the direction of endothermic change, i.e. towards the reactants, reducing the extra energy.

This follows Le Chatelier's principle.

The extra heat applied to the system is absorbed by the endothermic direction of change.


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Effect of temperature change on the equilibrium constant

Change of temperature actually changes the position of equilibrium. This means that it also changes the value of the equilibrium constant.

It is the only change of conditions that has this effect.

Change in conditions effect on reaction change Kc
concentration moves away from increased concentration no
pressure increased pressure moves away from greater moles no
temperature increased temperature moves towards endothermic change yes

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