 Hi students, welcome to year 12 chemistry and module 5 equilibrium and acid reactions. This is video number 7 looking at the effect of concentration on an equilibrium system. So in the last video we looked at the effect of temperature, very briefly looked at what would happen if we were to add or remove heat from the system and how the system would shift according to Le Chatelier's principle in order to counteract the changes that we made. In this video we're going to focus on the second of those important factors which is concentration. So here are two different systems that are in equilibrium. The first one is a heterogeneous system because we have a liquid as well as three species that are aqueous and the second of these systems is homogeneous because we're just considering the three ions in solution. So what we have whenever we look at an equilibrium system is we need to make sure that we have chemical reaction or an equation that represents that reaction in order for us to analyze what's going on. When we change the concentration there are a couple of things that we can do. We can directly add a particular salt or a compound to the solution that we know is going to make some sort of a change. The chromate ion here is the yellow ion so if we were to add for example potassium chromate so that would be a K2CRO4. So if we were to add this as just solid particles that we added to our mixture what we would find is that this would dissociate into K plus ions which are not going to actually interact at all but also into CRO4 to minus ions. Now this means that the change that we're making, this is the change, is going to increase the concentration of the chromate ions. This is the change. Le Chatelier's principle then says we need to do something to counter the change so the counter, the Le Chatelier's principle is a counter. Which means we need to drop the concentration. How do we drop the concentration? We drop the concentration by shifting the equilibrium to the right, shift, right. This favours the products and what it means is that the concentration after having increased will then start to drop for the chromate ions and it will increase for the dichromate ions. So we will shift from being a more yellowish colour into a more of an orange colour. So that's the macroscopic change you will see when we make these sorts of changes. One other thing that's useful to be aware of at this stage too is that sometimes what we do is we add something that looks unrelated, say something like sodium hydroxide. If we were to add a little sodium hydroxide to this particular reaction then we would have the sodium hydroxide dissociating into sodium ions and hydroxide ions. Now neither of these two ions are actually part of our equilibrium. However, one thing we know is that if we have hydroxide ions they will combine with hydrogen ions to form water. When that happens the concentration of the hydrogen ions will drop so effectively the sodium hydroxide is going to neutralise those H plus ions. So that again is our change. What we want to see happening is, so I'll just identify that, so there's a change. What we want to see is the counter to that change. So the counter to the change is wanting to increase the concentration of the hydrogen ions. In order to do that we're going to have to shift this equilibrium to the left. So this one's going to shift left to favour the formation of the reactants. Our orange colour is going to move more into the yellow range. This is what we do when we're looking at concentrations. You can see that the same sort of logic is going to apply if we were to add for example some iron 3 plus into our second equation. This is going to increase the concentration of the iron ions. This is our change. So what we need to do then is we need to look at our Le Chatelier's principle which is the counter. The counter is to try and get the concentration of the iron back down and therefore the only way to do that is to shift to the right so we use up those excess iron ions and we form more of the blood red ion thiocyanate ion as a product. So we shift to the right and we favour products. Carry out a few examples in order to reinforce that you are aware of exactly how changes in the concentration of different ions can actually shift the position of the equilibrium based on Le Chatelier's principle. Thanks for watching.