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 and reverse reactions are the same. We can show this as: reversible reaction 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 occur 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 above. The rate of the forward and reverse reactions can be altered if we change the reaction conditions, e.g. for example if we alter the concentration of one of the reactants or products or change the temperature of the equilibrium mixture, but recall that at equilibrium the rate of the forward reaction, Rf and the back reaction, Rb will be the same.

equilibrium 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 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 will re-adjust itself to get back to a new equilibrium position, that is a low energy state. At this new equilibrium position the amounts of reactants and products maybe 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 (reactants). Since ammonia is a useful substance we need to shift the position of equilibrium so that it moves more to the right (products).
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.
Le Chatelier's Principle is the idea that a system (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)   ΔH= -ve
Is exothermic, that is it releases heat energy to the surroundings. The back reaction:
2NH3(g)N2(g) + 3H2(g)   ΔH= +ve
Is endothermic, that is it removes heat energy from the surroundings.

ammonia, nitrogen and hydrogen in flask 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 position so that the amount of the valuable product ammonia in the equilibrium mixture 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:

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

To explain this observation we can simply use Le Chatelier's principle. The forward reaction:
Co(H2O)62+ + 4Cl-(aq) CoCl42- + 6H20   ΔH= +ve
Is endothermic, so if the violet solution is heated according to Le Chatelier's principle the equilibrium position will adjust itself to remove the added heat and establish a new equilibrium position. To remove the added heat the equilibrium position will shift to the right, that is turn the pink Co(H2O)62+ ion into the blue CoCl42- ion, so the solution obviously turns blue.
If we cool the violet solution containing the equilibrium mixture then what would you expect to happen? Again we can use Le Chatelier's principle to explain the colour change observed. If the forward reaction above is endothermic then the reverse reaction is exothermic. So according to Le Chatelier's principle if we remove heat from the system, then the equilibrium will try and oppose this change. To do this it will move the position of equilibrium to the left, that is produce more heat and also more of the Co(H2O)62+ ion, since this reaction is exothermic it will produce heat. So on cooling the colour of the solution will change from violet to pink.

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 an 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 equilibrium is affacted by changes in pressure

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 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.

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