When you think of chemical reactions, you might think of them as irreversible, permanently changing one substance into another. While this is true in some cases, some chemical reactions are reversible, and we can take the products of the reaction and turn them back into the reactants. These reversible reactions can, under certain conditions, reach what we call ‘equilibrium’. Equilibrium can be a tricky concept to understand, but this graphic tries to make it a little clearer.
Firstly, let’s clarify what equilibrium is. When we have a reversible reaction taking place in a closed system – that is, one where no substances are being added or lost – at the beginning of the reaction we will have only the reactants. Once the reaction starts, the amounts of the reactants will start decreasing, and the amounts of the products will begin to increase. In a non-reversible reaction this would be about the long and short of it, but when a reaction is reversible the products can also react to produce the reactants again.
After a time, a reversible reaction in a closed system can reach what we call a ‘dynamic equilibrium’. This is where the forwards reaction (reactants reacting to produce the products) and the backwards reaction (products reacting to reform the reactants) are occurring at the same rate. This means that the amounts of the reactants and the products remain constant, despite the fact that both reactions are still ongoing.
As an analogy for this, it’s helpful to think of two men, one shovelling dirt out of a hole, and the other standing outside the hole shovelling dirt back in. If both are shovelling at the same rate, then the size of the hole and the size of the pile of dirt outside the hole will remain constant, despite the fact that both of them are still shovelling dirt back and forth. Dynamic equilibrium is much the same.
Though the amounts of the reactants and products may initially be the same once we reach dynamic equilibrium, we can make adjustments to the conditions of the equilibrium to change the proportions of reactants and products in the equilibrium mixture. The results of changes we make can be determined using something called ‘Le Châtelier’s Principle’.
Le Châtelier’s Principle states that when we make changes to a reaction at equilibrium, the equilibrium will respond to the change we make to try and undo the change. For example, if we increase the temperature of the reaction, it will respond in a way that decreases the temperature. If we increase the concentration of a reactant, it will respond in a way that decreases the concentration of that reactant. On the face of it, this seems quite simple. However, applying it can be a little trickier than it initially seems.
Let’s start with discussing concentration. Concentration is simply a measure of how much of a substance we have in a particular volume. If we increase the concentration of a particular substance in the reaction, according to Le Châtelier’s Principle the equilibrium’s response will be to try and reduce the concentration of this substance.
Say we increase the concentration of one of the reactants; the equilibrium can reduce its concentration by favouring the forwards reaction and producing more of the products. We sometimes phrase this as the equilibrium ‘moving to the right’. The net result of increasing the concentration of the reactants would be the production of more of the products at equilibrium.
Changing the temperature can also affect equilibrium position. If we increase the temperature, according to Le Châtelier’s Principle the equilibrium will act to reduce the temperature. How it does this and whether it favours the reactants or the products will depend on the reaction.
Chemical reactions can be either exothermic (give out heat) or endothermic (take in heat). If the energy required to break bonds is less than the energy released when forming new bonds, the reaction will be exothermic. If the energy required to break bonds is more than the energy released when forming new bonds, the reaction is endothermic. For reversible reactions, either the forwards or backwards reaction will be exothermic, and the other will be endothermic.
When we increase the temperature, the reaction will favour whichever reaction is endothermic to take in heat and reduce the temperature. On the other hand, if we decrease the temperature, the exothermic reaction will be favoured, as this will give out heat and increase the temperature.
Changing the pressure of a reaction involving gases can also affect the position of equilibrium. Pressure is caused by the collisions of gas particles with the walls of the container. It follows, then, that the greater the number of gas particles, the higher the pressure will be – and this gives us a hint as to the effect changing pressure has on the equilibrium position.
If we increase the pressure, Le Châtelier’s Principle states that the equilibrium will change to reduce the pressure. It can do this by favouring the side of the reaction with fewer gas molecules; which side this is will of course depend on the reaction in question. Conversely, if we decrease the pressure, the equilibrium will respond to increase the pressure, and will therefore favour the side of the reaction with more gas molecules.
As a caveat, there are plenty of reactions where we actually have the same number of gaseous molecules shown in the balanced equation. In this case, changing the pressure will have no effect, as it will not favour either side of the reaction.
Catalysts are often used by chemists to increase the speed of chemical reactions. Using one doesn’t affect the position of equilibrium in any way, however. This is because it speeds up both the forwards and backwards reaction by an equal amount, so overall neither side of the reaction is favoured by using one.
Why does this all matter?
You could be forgiven for wondering why chemists are so interested in how changing the conditions of an equilibrium will affect it, and why it really matters to anyone who’s not a chemist. It’s actually more important for non-chemists than you might think. The Haber process, a reaction in which hydrogen and nitrogen are combined to make ammonia, is an example of a reaction which takes advantage of some of the chemistry we’ve discussed here. Some of the ammonia produced by the process is used to make fertilisers, and it’s estimated that this helps feed about 40% of the world’s population.
The problem with the reaction is that it is reversible; therefore some tweaking is required to produce as much ammonia as possible. Firstly, during the process any nitrogen and hydrogen that doesn’t react is recycled until it does. The ammonia that is produced is also removed from the reactor, so that the equilibrium responds by shifting to produce more ammonia. A iron catalyst is often used to make the reaction happen faster.
In terms of temperature, the forwards reaction (nitrogen and hydrogen to ammonia) is exothermic. This means that it’s favoured by using a low temperature. However, industrially the Haber process is usually carried out at 400-450˚C. This is by no means a low temperature, so why is it used? Well, temperature also affects the rate of the reaction, as detailed in a previous graphic here – if the temperature is too low, the reaction will happen too slowly! Therefore, what’s called a ‘compromise temperature’ is used, striking a balance between a fast enough reaction and conditions which favour the formation of the desired products.
Finally, the balanced equation for the reaction shows that there are more molecules of gas on the left-hand side of the equation (4) than there are on the right-hand side (2).
Consequently, the reaction is favoured by a high pressure, as this shift the equilibrium towards the side with fewer molecules of gas, producing more ammonia. For this reason the Haber process is usually carried out at pressures around 200 atmospheres. Again, there is a degree of compromise here; higher pressures would produce even more ammonia, but would require more expensive equipment and be more costly to maintain.
As a final word, it’s worth pointing out that there are other ways of considering how changes affect equilibria by using more mathematical approaches instead of Le Châtelier’s Principle. These are better at aiding predictions, as they can avoid some of the confusion that the Principle can sometimes result in. However, if you’re a student here in the UK these methods don’t appear until after GCSE, so Le Châtelier’s Principle is, like it or loathe it, your only tool for predicting equilibrium until then!
Enjoyed this post & graphic? Consider supporting Compound Interest on Patreon, and get previews of upcoming posts & more!
References & Further Reading
- Le Châtelier’s Principle – J Clark, Chemguide
- A rant against the use of Le Châtelier’s Principle – E Scerri
- The Haber Process for the manufacture of ammonia – J Clark, Chemguide