I’m sure you will all have seen a strip of magnesium burning in a Bunsen burner flame. It is a highly exothermic reaction producing lots of white hot sparks, and at the end we are left with a white powder of magnesium oxide.
Mg(s) + ½O2(g) ⇾ MgO(s)
This reaction goes to completion, and however long you stare at the white powder, it never unburns to reform the magnesium strip!
However, many physical and chemical reactions don’t go to completion – they reach a state of balance or a state of equilibrium in which some of the reactants have formed products, but not all. After a certain period of time the reaction mixture contains both reactants and products and the amounts of each doesn’t vary noticeably.
Consider a wet sock sealed in a plastic bag. We have a closed system. Neither the water nor the sock can escape from the sealed bag, which means that the sock never dries out.
- water evaporates from the sock into the bag’s atmosphere
H2O(l) ⇾ H2O(g)
- eventually the atmosphere in the bag becomes saturated and water begins to condense back onto the sock
H2O(g) ⇾ H2O(l)
- after some time an equilibrium is established, with water evaporating and condensing at the same rate so the sock never gets any wetter or any drier
H2O(l) ⇌ H2O(g)
This is a dynamic equilibrium because both the forward and back reactions (the evaporating and the condensing) continue to happen, even though we don’t notice.
We can illustrate this graphically.
Once equilibrium has been established the rate of the forward reaction equals the rate of the back reaction. The amounts of the reactants and products don’t change (they are constant) but that does not mean that they are necessarily the same. The actual amounts of gaseous and liquid water in our bag at equilibrium depend on how wet the sock was to begin with, the humidity of the air, the temperature and the pressure in this closed system.
If this was chemical reaction rather than a physical one, the ‘amounts’ of reactants and products present in the system at equilibrium would be the concentration of each if we were working with solutions or the partial pressures of each if we were working with gases.
We use the term position of equilibrium to describe one particular set of equilibrium concentrations for the reactants and the products in a system.
So what exactly do we mean by that?
Let’s place the bag in a greenhouse on a hot, sunny day. The temperature in the bag (closed system) rises.
We would expect that the rate of evaporation will increase initially as at a higher temperature the air can hold more gaseous water, but after a certain period of time the system will establish a new equilibrium. Once again the rate of evaporation and condensation are equal when our system is in equilibrium, but the position of equilibrium has changed.
H2O(l) ⇌ H2O(g)
The position of equilibrium has shifted to the right (in the context of the equation as we have written it) because at equilibrium we have a higher amount of gaseous water than we had previously.
Similarly, if we put the bag in the fridge, less water will evaporate. When the dynamic equilibrium is established inside the bag we would expect the new position of equilibrium to reflect a much higher amount of liquid water and far less gaseous water than before.
Language is absolutely key in this topic, I can’t stress this enough! The video reinforces how to explain these concepts (and how not to) in exam speak👇.
I’d also recommend making some flashcards for revision with clear definitions of all the key terms.