 One of the things that's very important as a consequence of the Bronsted-Larrow definition of acids and bases is that we can identify conjugate pairs. Now conjugate pairs are acid-base pairs that just differ by one proton in their formula. There's a huge number of examples of these and we'll look at all of these in a range of different reactions, but you can see just the one that I've put here, the ammonia, NH3 form, and the ammonium iron, NH4 plus form differ in the presence or absence of a proton, an H plus, so if I add an H plus to NH3 I get NH4 plus, or if I subtract NH plus from NH4 plus, I get NH3. The same is true of hydrochloric acid and the chloride iron. It's true of water and the hydronium iron. It's also true of water and the hydroxide iron, and there are an endless number of examples of these. All of these are conjugate pairs. One is an acid and one is a conjugate base. Sometimes you might describe one as a base and it's conjugate acid, but the difference between these is simply that proton, that H plus that has either been donated or accepted by one or other of these species. One of the important things about conjugate pairs is that we can actually use them to start to get a sense of the strength of different types of acids or bases and this is, I guess, independent of what we've looked at previously, which were mathematical expressions for the difference between a strong and a weak acid. Now we get a chance to have a look at what's happening in terms of the conjugates. The equilibrium position of the acid-base system is actually, generally, determined by the relative strength of the base. So if we know what the base is in a conjugate pair, then we can get a bit of a sense of the strength of its conjugate acid. So if, for example, I pick the top one, if the base is the chloride ion, then the conjugate acid, if I add an H and a plus, I get HCl. The fact that this is up there on the strongest of the acids or, more specifically, on the weakest of the bases. So if we are trying to analyze base strength here, then chlorine, the chloride ion, is actually at the lowest end of that scale. It is the weakest base in terms of strength. Therefore, it's going to form, if this is the weakest base, it's going to form the strongest acid. And this is one of the important things about our conjugate pairs. The pair strengths are the reverse. A strong acid has a weak conjugate base and a strong base has a weak conjugate acid. Where we look at something that we may be a little bit familiar with somewhere in the middle of the table here. This is acetic acid. It's the acid that's present in vinegar, also known as ethanoic acid. It's kind of middle of the table. It's conjugate base here. And you can see, again, the difference between these two is just an H plus. But because this is kind of a mid strength middle of the table in terms of its base strength, then we know that the conjugate acid, the corresponding acid for the acetate ion, is also going to be relatively weak. Certainly not as strong as hydrochloric, but stronger than some other options. The one directly underneath it with the bicarbonate ion, HCO3 minus. This is the bicarbonate ion. And so therefore, its conjugate acid would be carbonic acid. And again, we know this is the acid that's present in carbonated soft drinks. So again, it's a weak acid. It's one that we can consume. And again, it's it's kind of middle of the table. You have to kind of go almost to the bottom to see something that's really weak in terms of acid strength. And the one that's second from the bottom is one that kind of probably makes a little bit of sense to us, the hydroxide ion. When the hydroxide ion is in solution, it's very, very strong. So it does not readily accept a proton to form water. And as a consequence of that, we find a lot of hydroxide ions are not drawing hydrogen ions out of solution to form water molecules. We find those hydroxide ions can dominate in a solution. So a strong base has a weak conjugate and a weak base has a strong conjugate.