Header image Le Chatelier's principle.

Le Chatelier's principle

Henry Le Chatelier was a French chemist who studied reversible reactions and equilibrium and he made some important observations.

Equilibrium is the point in a reversible reaction where the rate of the forward (Rf) reaction and reverse reaction (Rb) are the same. We can show this as: An equation to show that a reaction is reversible. During a reversible reaction the reactants turn into products and the products back into reactants again. The important point to remember is that these two reactions are occurring at the same time. The forward reaction will proceed at a given rate; labelled Rf and the back reaction which turns products back into reactants will also proceed at a given rate; labelled Rb in the example below.

Equation for a reaction at equilbrium.

When a chemical reaction starts usually the rate of the forward reaction is very fast; this is simply because there are lots of reactant particles around to collide and react with each other. However as the reaction proceeds the rate of the forward reaction (Rf) will start to slow down as the amount of reactants particles available to react reduces. The opposite is true for the back reaction; to begin with there will not be many products particles available to react with each other but as the reaction proceeds the amounts of products present will gradually increase and the rate of the back reaction (Rb) will increase. Eventually after a certain period of time the rate of the forward and reverse reactions will be the same.

Imagine looking a reversible chemical reaction where the rates of the forward and reverse reactions were the same; what do you think you would see happening? Well as quickly as the reactants are turning into products the products are turning back into reactants, so the amounts of the reactants and products will not be changing. The reaction might appear as if it has stopped but it has NOT. The reactants are turning into products and the products are turning back into reactants at the same rate; the reaction most definitely has not stopped; it has reached a balance point where the rates of the forward and reverse reactions are the same. As a chemist we would say the reaction has reached or achieved dynamic equilibrium. The word equilibrium implies balance and dynamic implies movement. The reaction has reached equilibrium because Rf=Rb but the reaction is still occurring, its dynamic!

Altering the position of equilibrium

Model to explain equilibrium.

From the work you have done on rates of reaction you probably already know that it is possible to change the rate of a reaction by say heating it up or increasing the concentration of a particular reactant, adding a catalyst or increasing the surface area of a reacting substance. A common misconception that a lot of students have is that if a reaction has achieved dynamic equilibrium then it consists of a 50:50 mixture of reactants and products; however this is rarely the case.

Equilibrium is a very stable low energy point for a reaction; at equilibrium the reaction is nice and happy!! It's like the ball at the bottom of the curve. If you push on it so that it moves away from the bottom of the curve the ball will roll back down again; to the low energy point. Well chemical reactions are a bit like that!- if a reaction is at equilibrium (rate of forward and reverse reactions are the same) and you come along and heat it up or put it under pressure the reaction and generally annoy it then the reaction will re-adjust itself to get back to a new equilibrium position; that is a new low energy state. At this new equilibrium position the amounts of reactants and products may be different but the rate of the forward and reverse reactions will be the same e.g. consider the reaction of nitrogen and hydrogen to make ammonia; the Haber process.

Nitrogen(g) + hydrogen(g) ammonia (g)
N2(g) + 3H2(g) 2NH3(g)

At equilibrium the mixture contains only a small amount of ammonia and is mostly nitrogen and hydrogen. We would say the position of equilibrium lies to the left; that is on the reactants side. Since ammonia is a useful substance we need to shift the position of equilibrium so that it moves more to the right, that is the equilibrium mixture has a larger percentage of ammonia; the product present.

The position of equilibrium for a reaction depends on several variables; these include concentration, temperature and pressure if gases are involved in the reaction. We can adjust the amounts of reactants and products at equilibrium by simply changing the reaction conditions. The trick is to know exactly what conditions to change and that is exactly what Henri Le Chatelier figured out.

In chemistry we often use the words system and surroundings when discussing a particular reaction. The system is simply the reacting chemicals and the surroundings are the test tubes, beakers and indeed the whole universe- it's everything except the reacting chemicals! A closed system is one where no reactants or products can escape from the apparatus and nothing from the surroundings can get in either. Le Chatelier's Principle is the idea that a system (the reacting chemicals) at equilibrium will oppose any changes applied to it. We will use the Haber process to try and explain Le Chatelier's principle in more detail. The equation for the Haber process is shown again below:

Nitrogen(g) + hydrogen(g) ammonia (g)
N2(g) + 3H2(g) 2NH3(g)
The forward reaction:
N2(g) + 3H2(g) 2NH3(g)
is exothermic; that is it releases heat energy to the surroundings. The back reaction:
2NH3(g) N2(g) + 3H2(g)
is endothermic; that is it removes heat energy from the surroundings.

ammonia, nitrogen and hydrogen in flask to demonstrate how changes in temperature, pressure and concetration of gases will change the position of equilibrium.

So if you have a mixture of nitrogen, hydrogen and ammonia in a flask at equilibrium. The position of equilibrium for this reaction lies very much to the left and unfortunately there is very little ammonia in the flask. The challenge for chemists is how to adjust this equilibrium mixture so that the amount of the valuable product ammonia can be increased.

Example 2 - Cobalt (II) Chloride Temperature and Le Chatelier's principle

Cobalt (II) choride (CoCl2) will dissolve in water to form the complex ion Co(H2O)62+ which can be dissolved endothermically in dilute hydrochloric acid to form the deep blue complex CoCl42-. This reaction which is shown below is endothermic.

Co(H2O)62+ + 4Cl-(aq) CoCl42- + 6H20   ΔH= +ve
Since this is a reversible reaction it will contain an equilibrium mixture of the pink Co2+ ion and the blue CoCl4- ion and this mixture is a lovely violet colour, see image below:

The effect of heat on an equilibrium mixture of coloured transition metal compounds.

As shown in the image above when the violet solution is:

To explain this observation we can simply use Le Chatelier's principle.

Example 3 - Nitrogen dioxide and nitrogen tetraoxide. Pressure and Le Chatelier's principle

Equation showing the equilibrium reaction between nitrogen dioxide and nitrogen tetraoxide gas Using Le Chatelier's principle to explain changes in the position of equilibrium due to changes in pressure only applies to reactions involving gases. Solids and liquids are not affected by changes in pressure. As an example consider the reversible reaction that occurs between the colourless gas nitrogen tetraoxide (N2O4) and the brown gas nitrogen dioxide (NO2).

You can see from the equation opposite that 1 mole of the colourless N2O4 gas dissociates to form 2 moles of the brown NO2 gas. Since there is 1 mole of gas on one side of the equation and 2 moles on the other we can say that the side with 1 mole of gas will be the low pressure side and the side with 2 moles of gas will be the high pressure side. The image below shows what happens to an equilibrium mixture of N2O4 and NO2 gases in a syringe as the pressure is reduced and increased: how the position of equilibrium is altered by changes in pressure for an equilibrium mixture of nitrogen dioxide and nitrogen tetraoxide.

Example 4 - Concentration and Le Chatelier's principle

Consider the following reversible reaction which involves only solutions.

A(aq) + B(aq) ⇌ C(aq) + D(aq)
Again we can use Le Chatelier's principle to figure out what will happen to the concentration of reactants and products as we add or remove any of them. Suppose for example we add more of substance A to the equilibrium mixture, what will happen to the concentration of the other reagents?

In a closed system adding more of A will increase its concentration; so according to Le Chatelier's principle the equilibrium mixture will try to oppose this change; that is reduce the concentration of A. The only way to do this is to react B with A to reduce its concentration, this will make more of the products B and C and will obviously reduce the amount of A and B present. The position of equilibrium will move to the right.

The same affect would happen if we removed or reduced the concentration of the product C; this time according to Le Chatelier's principle the system will try to make more C by reacting together A and B. By continually removing C the reaction will never be able to achieve equilibrium. For reactions such as the Haber process discussed above by continually removing ammonia from the equilibrium mixture we are in effect forcing more hydrogen and nitrogen to react.

Key Points

Practice questions

Check your understanding - Questions equilibrium conditions.