Balancing stones

What is an Equilibrium Reaction?

If you’re studying HSC Chemistry, you’ve probably started to encounter equilibrium reactions. From our tutoring courses in the summer holidays, I realised that this can be a challenging topic for many students. Hopefully this article will clear up misconceptions and cement your understanding of the concept.

In your study of chemistry so far, you’ve seen many different chemical reactions take place. However, in each case the reactants convert entirely into the products. You’ve probably used stoichiometry to determine the exact yield of a reaction given a known amount of the reactants (and you’ll continue to do that this year so make sure to stay on top of the calculations). This is only possible because we know that the reaction ‘goes to completion’ and no reactants remain.

However, this is not true of all reactions. A significant number of chemical reactions are ‘equilibrium reactions’, which can also be referred to as ‘reversible’ or ‘non-spontaneous’. In these reactions, the forward reaction does not go ‘to completion’; instead, a certain amount of the reactants remains.

At a molecular level, the reverse direction of the reaction (from products to reactants) is also occurring. This means that the reactants can never be fully consumed (they are being recreated by the reverse reaction). More importantly, if the rate of the forward reaction and reverse reaction are not the same (a likely scenario), then the amounts of the reactants and products will be changing. When the reaction rates become the same, then the amounts will become constant. At this point, we say that the system is in ‘equilibrium’. More specifically, we call it a ‘dynamic equilibrium’ because while the total amounts of each stay the same, individual molecules are still being converted between the reactants and products.

But how do we know that the reaction rates will ever be the same? Remembering back to year 10 science (or even earlier), the rate of a reaction is proportional to the concentration of the reactants. That means that if more of the reactants are present, the reaction will occur more quickly. The result in our equilibrium reaction is that as the reaction proceeds in the forwards direction, the reactants will be used up and the rate of the forward reaction will reduce. Now consider the rate of the reverse reaction: it is proportional to the concentration of the ‘products’ of the original reaction. Therefore, as this reaction proceeds and more products are produced, the rate of the reverse reaction will increase. We can visualise this process with the graph below to realise that an equilibrium point will always be established if the reverse reaction is possible.

via edublognss.wordpress.com

However, the location of that equilibrium point depends on the particular reaction. In some cases, we say that an equilibrium ‘lies to the right’, which means that there are more of the products than the reactants at equilibrium. In contrast, an equilibrium which ‘lies to the left’ has more of the reactants than the products at the equilibrium point. In these cases the reaction will only proceed a small amount before equilibrium is reached.

Finally, let’s go through an example:

H2O (l) + CO2 (g) ⇌ H2CO­3 (aq)

The first thing to notice is the different arrow we use. The bi-directional arrow is used in chemical reactions to indicate that an equilibrium is established. Under standard conditions (1 atm of pressure, 25 degrees Celcius), this equilibrium lies to the left. This means that only a small amount of the CO2 present in a system will be dissolved in water. Once we reach an equilibrium, both the forward reaction, where CO2 dissolves in water, and the reverse reaction, where it leaves the solution are occurring. However, the rate of these reactions is the same, so the relative amounts of dissolved and undissolved CO2 will not change.

I hope this article clears up some misconceptions and/or assists your understanding of equilibrium reactions. In my next article, I’ll explain changes that occur in equilibria through Le Chatelier’s principle. To make sure you don’t miss out, sign up for our newsletter at the top right of this page and get all of our articles straight to your inbox.

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