When the forward and the backward reaction reach the same speed, there is dynamic equilibrium. The quantities of the reactants and products will not

change unless they are forced to.

When dynamic equilibrium is

reached, there does not have to be equal amounts of reactants and products in the

equilibrium mixture. If there are more reactants than

products we say the equilibrium position

lies to the left. If there are more

products then we say the equilibrium

position lies to the right.

This states:

The position of equilibrium shifts to try to cancel out any changes you introduce.

For example if you add more reactants the equilibrium position will change in order to “get rid of” the extra you have added. This means you will get more products. If you remove the products as they are made, then the equilibrium position will shift to

the right to compensate.

The equilibrium position can also be changed by altering temperature and pressure.

Altering Pressure (an example)

N2 + 3H2 2NH3

High pressure as there are 4

molecules hitting the sides of the

container

Low pressure as there are 2

molecules hitting the sides of the

container

If the pressure is increased the equilibrium position will move to the right. If the pressure is decreased the equilibrium position will move to the left. Why?

Altering Temperature (an example)

H for this reaction is exothermic. This tells us that the forward reaction is

exothermic and the backward reaction is endothermic.

N2 + 3H2 2NH3

EXOTHERMIC

ENDOTHERMIC

If the temperature is increased the equilibrium position will move to the left. If the temperature is decreased the equilibrium position will move to the right. Why?

Fritz Haber made an equilibrium mixture containing nitrogen,

hydrogen and ammonia in 1909. However, with the

equipment available he could make only 100g of ammonia.

The German company BASF tried to scale up

the process. The chemical engineer Carl

Bosch solved the problem of needing high pressure. He devised a double-walled steel vessel which would work

safely at the 300 times atmospheric pressure.

The production of ammonia is called the Haber Process. The conditions required for the process have been carefully selected to get the best yield of ammonia in the quickest time possible, making it financially viable.

The conditions are:

• Pressure between 150 - 300 atmospheres

• Temperature between 400 - 450°C

• Use of an iron catalyst

• Unused nitrogen and hydrogen are recycled

• The ammonia is removed as it is made

The first graph shows as the pressure is increased the yield of ammonia goes up (at constant temperature). However, above 300 atmospheres of pressure it becomes more costly and more dangerous.

The second graph shows as the temperature is increased the yield of ammonia goes down (at constant pressure). However, higher temperatures are used so that you make the ammonia more quickly.

The iron catalyst does not affect how much ammonia you make - only how quickly you get it!