Henry Le Chatelier was a pioneering French chemist born in 1850, whose work and research have enable scientists to better understand chemical equilibrium. His most famous contribution is perhaps the principle that bears his name, Le Chatelier's Principle. This is a theory to help chemists and other scientists understanding how chemical systems (the reacting chemicals) at equilibrium respond to changes in the reaction conditions, whether it be changes in the temperature, pressure, or concentration of a reactant or product. In essence, this principle states that a chemical system will change its composition to minimise the effect of any changes in temperature, pressure or concentration that are made to any of the substances involved in the equilibrium reaction.

Equilibrium is the point in a reversible reaction where the rate of the forward (R_f) reaction and reverse reactions (R_b) are the same. We can show this as: 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 R_f and the back reaction which turns products back into reactants will also proceed at a given rate; labelled R_b in the example below.

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 (R_f) 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 (R_b) 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 R_f=R_b but the reaction is still occurring, its dynamic!

Altering the position of 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)

N_2(g) + 3H_2(g) ⇌ 2NH_3(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 in order to change the position of equilibrium and the amounts of the reactants and products present 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)

N_2(g) + 3H_2(g) ⇌ 2NH_3(g)

N_2(g) + 3H_2(g) → 2NH_3(g)

2NH_3(g) → N_2(g) + 3H_2(g)

So if we have a mixture of nitrogen, hydrogen and ammonia gases at equilibrium in a flask at room temperature and pressure, unfortunately the position of equilibrium for this reaction lies very much to the left and 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, so:

• What would happen if we warmed the flask? Well according to Le Chatelier's principle the reaction should oppose any change you make to it; that is it will try and remove the heat you added. How could the added heat be removed? An endothermic reaction will remove heat from the system. The only way to do this is to force the equilibrium position even more to the left, that is to increase the back reaction since this is an endothermic reaction:

2NH_3(g) → N_2(g) + 3H_2(g)

• However this is bad news as it will reduce the amount of ammonia in the flask. So to increase the amount of ammonia instead of heating the flask, cool the mixture in the flask. Again according to Le Chatelier's Principle the system will try to generate heat; to oppose the cooling; so more nitrogen and hydrogen will react to make ammonia since the forward reaction is exothermic. However we know that cooling a reaction will slow it down, it will reduce the rates of the forward (R_f) and back reaction (R_b). So ultimately cooling it down will increase the amount of ammonia in the flask but it might take a very long time for the new equilibrium position in which there is more ammonia present to become established.
• What would happen if you increased the pressure to the position of equilibrium in the Haber process? Well first thing is that pressure will only have an effect if there are gases involved in the reaction. So since all the chemical reacting here are gases then pressure will have an effect. At room temperature and pressure one mole of any gas will occupy 24 litres. So for the reactants the total number of moles is 4 so this will occupy a volume of 96 litres. For the products we have 2 moles of gas; so this will occupy 48 litres at room temperature and normal pressure (standard temperature and pressure S.T.P). So obviously we can think of the reactants as the high pressure side of the equation and the products as the low pressure side. So if you increase the pressure on this equilibrium mixture of gases then it will force the equilibrium to the right, that is the equilibrium mixture will contain more ammonia. If you reduce the pressure then the equilibrium mixture will contain more reactants and less ammonia.
• Similar arguments are true if you add more reactants or products then the system will no longer be at equilibrium and it, the system will adjust the concentrations of all reactants and products to achieve a new equilibrium position. If you add more reactants it will force the equilibrium to the right, that is towards the products, similarly if you add more products the system will readjust to try and remove the added products by pushing the equilibrium towards the left hand side, that is the amount of reactants will increase.
• How about adding a catalyst? Catalysts do not have any effect on the position of equilibrium; they will not affect the amount of reactants or products in a equilibrium mixture. What a catalyst will do is increase the rate of both the forward and reverse reactions in an equilibrium mixture. The catalyst will allow the system to achieve equilibrium faster.

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 CoCl₄^2-. This reaction which is shown below is endothermic.

Co(H₂O)₆²⁺ + 4Cl^-_(aq) ⇌ CoCl₄^2- + 6H₂0   ΔH= +ve

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

• Heated it turn blue
• When cooled it turns pink

• The forward reaction:

  Co(H₂O)₆²⁺ + 4Cl^-_(aq) ⇌ CoCl₄^2- + 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 in the direction of the endothermic reaction. That is turn the pink Co(H₂O)₆²⁺ ion into the blue CoCl₄^2- 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

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:

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)

Key Points

• Dynamic equilibrium is the point in a reversible reaction where the rate at which the reactants form products is exactly matched by the rate at which products reform reactants.
• It is only possible to achieve dynamic equilibrium in reactions which are closed. This means no reactants or products are allowed to escape or leave the reacting system, e.g. no gases are allowed to escape otherwise equilibrium will never be achieved.
• Le Chatelier's principle states that a system at equilibrium will oppose any changes made to it.
• Changing the temperature and concentration of the reactants or products will effect the position of equilibrium in a reversible reaction. Changing the pressure of a reacting system will only alter the position of equilibrium if gases are involved in the reaction.
• Catalysts DO NOT change the position of an equilibrium mixture, this simply menas that a catalyst will not alter the amounts of reactants or products in a reversible reaction that is at equilibrium. A catalyst will enable the reaction to get to equilibrium much faster.

Check your understanding - Questions equilibrium conditions.

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