Many chemical reactions are reversible. In these reactions, there is both a forward reaction (where reactants are made into products) and a reverse reaction (where product molecules break down to form reactants).
The Haber process, the industrial route to the formation of ammonia from nitrogen and hydrogen gas, is an example of a reversible reaction.
In a reversible reaction, you can never obtain 100 per cent conversion of reactants into products. Reversible reactions will always result in a mixture of reactants and products being formed.
While this isn't a major problem with the Haber process, it can often result in expensive reactant molecules not being completely converted into products.
As the forward reaction slows down, the reverse reaction will speed up until they are both taking place at the same rate. This is called the equilibrium position.
At equilibrium the concentration of reactant and products remain constant but NOT necessarily equal.
Equilibrium can only be obtained in a closed system where the reaction is carried out in a sealed container and none of the reactants or products are lost.
For example, iodine crystals break down to form purple iodine vapour. In an open system, the vapour escapes and the reaction progresses until all of the crystals have vaporised, and all of the vapour has escaped.
If a stopper is placed on the boiling tube, a closed system is formed. The iodine crystals break down to form purple iodine vapour, but both the crystals and vapour remain.
In the closed system, equilibrium has been established.
In an open system, products (or reactants) are lost, therefore equilibrium cannot be established.