An Introduction to Equilibria

Equilibria can be a tricky concept to understand. In order to understand it, you need a good idea of reversible reactions and then a good idea of a dynamic equilibria.

So, what is a reversible reaction?

Put simply, it is a reaction in which products can be turned into reactants, but reactants can also be turned back into products.

This can be shown using the ⇌ symbol:

Reactants ⇌ Products

For example, ammonium chloride will thermally decompose into ammonia and hydrogen chloride gases, but these gases will also recombine into ammonium chloride.

This is shown as follows:

NH4Cl (s) ⇌ NH3 (g) + HCl (g)

Notice how the ⇌ sign indicates a reversible reaction.

A reaction that is in dynamic equilibria takes this concept a step further. In this sort of reaction, there will always be reactants and products (the reaction never ends) and the proportion of reactants and products will stay constant (unless an external change is applied to the reaction).

A classic example of this is the production of ammonia in the Haber process:

3H2 (g) + N2 (g) ⇌ 2NH3 (g)

Even though the proportion of reactants and products stays constant, the reaction is still occurring. How can this be?

Well, what is happening is the rate of the reaction going from left to right (the ‘forwards’ reaction) is occurring at the same rate as the reaction going from right to left (the ‘backwards’ reaction).

Reactants are still reacting to form products but products are also decomposing to reactants at the same rate—the reaction is still occurring—it is dynamic.

Part of the trick in industrial reactions that are at equilibrium is to ‘shift’ the equilibrium to the product side (to the left) by applying some external change to the reaction.

In order to predict what happens, we need to understand Le Chatelier’s principle (or, as I put it, the law of opposites).

This principle works on the fact that whatever you do, the reaction does the opposite.

You increase the concentration of the reactants—the reaction tries to decrease them.

You decrease the pressure—the reaction tries to increase them.

You increase the temperature—the reaction tries to decrease the temperature.

Get the idea?

In order for the reaction to do the opposite to what you have done, it will shift the equilibrium away from the change. So, if you increase the concentration of reactants, the equilibrium will shift it away from this change, to the product side. The effect of this is the change is reduced and more products formed.

The following video clip will help you to get to grips with some of these concepts:

Source: © 2013 FuseSchool
Available under a Creative Commons license (CC BY-NC-ND 3.0), via YouTube

Which part of the equilibrium topic do you find most challenging? Post your comments and questions below and I will do my best to answer them!

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