 So in this particular video, we're going to put a few things together just to see if we can understand what buffers are and how we can use them. And therefore, we can then put together a practical in class about buffer and just demonstrate how it acts as a buffer. So firstly, a couple of definitions. One from Chemistry in the Marketplace by Selinger and Barrow, and this is a great book. A buffer is a chemical added to water to minimise the change in the pH if acids or bases are added. Bronwyn Hagerty's work, her module on acid-base reactions also has a nice definition. It's a longer definition and I've just kind of trimmed out the key bits here for us to focus on. Buffer solution is an aqueous solution consisting of a mixture of a weak acid in its conjugate base or a weak base in its conjugate acid. So here we have a couple of definitions that give us a bit of an idea about what's going on with a buffer. We can use both our knowledge of conjugates and also an understanding of our equilibrium constant in relation to the acid ionisation to give us a little bit more of an understanding of what's going on here. So if, for example, we were to look at a weak acid and its conjugate, so let me just pick something like acetic acid. So here's acetic acid. Now we know that acetic acid in water will act as a Bronsted Larry acid and it will donate a proton to the water. So we'll get an H3O plus and we'll also have the CH3C00-acetate ion. So this is a standard. Here is our weak acid and here is our conjugate base. So we've identified both of these two things. But what we know is that because this is weak, there is an equilibrium that lies to the left here. So it favours the formation or at least the retention of the molecule, not the ion. So how do we make a buffer in the first place? Well what we have to do is we have to add a common ion. So for example, if we added sodium acetate, then what we would do is increase the concentration of the acetate ion. Now of course, we know that in equilibrium systems there's going to be a shift to the left to try and re-establish the equilibrium. But if we add sufficient conjugate base to our solution, then we will effectively have equal concentrations of both our acid and our conjugate base. What we want is for these to be equal. Now the thing that's really important when these two things are equal is that when we look at the Ka value, which is the concentration of in this case the H3O+, multiplied by the concentration of CH3C00- divided by the concentration of the CH3C00H. When we look at this particular expression, we notice that if the concentration of the acid and the conjugate are the same, they will cancel out and therefore our Ka value will effectively be tied directly to the concentration of hydrogen ions. And that's what we're after. We're actually after a particular type of buffer that has a value that's associated with that concentration of hydrogen ions. This means that there are different types of bases and the Hegarty definition makes it clear. There are different types of bases of conjugate acid or base combinations to form buffers that will have different pH values. So we can keep the buffer at a low pH, at a neutral pH or at a higher pH depending on which substance we use in combination with its conjugate. The most important thing for us therefore to do when we're looking at producing a buffer is to ensure that we have both the acid present and its conjugate base in the form of a salt usually, or vice versa if we're starting with the base.