Equilibria

There are quite a number of reversible reactions, for example, the manufacture of ammonia. During a reversible reaction, both the forwards reaction of reactants going to products AND the backwards reaction of products going to reactants are taking place at the same time. At equilibrium, the two chemical reactions are going at the same rate and so the concentrations of each substance present in the mixture remains the same.

It is possible to disturb equilibria in reversible reactions by changing the conditions. If you understand rates of reaction, understanding equilibria should be reasonably straightforward for you. In industrial chemistry, getting the balance right is often a compromise between the speed of reaction and the yield, as you will have seen in your studies of the Haber Process.

Firstly, let's look at an increase in temperature. In reversible reactions, if one of the reactions is exothermic and the other endothermic, the yield of products will be affected. A temperature increase will favor the endothermic reaction since there will be more energy available to be absorbed from the surroundings. So if the forward reaction is endothermic, a higher percentage of the reactants will be converted into products, and vice-versa if it is the reverse reaction that is endothermic.

Secondly, we will consider what effect an increase in concentration has on equilibria. If the concentration of the reactants is increased, that will increase the rate of the forward reaction, so more product will be made and vice-versa. Where the reaction is between gasses, the equivalent of an increase in concentration is an increase in pressure. If there are fewer gas molecules on the products side than on the reactants side, the higher pressure will favor the products. As always, with equilibria, the opposite is also true.

When discussing changes to equilibria, drugstores usually talk about pushing the equilibrium to the left or to the right. An equilibrium that is pushed to the right actually indicates that the forward reaction is favored and that you will get more product; one that is pushed to the left will have less product and more reactants.