Buffers

Okay, a review of buffers. The definition of a buffer is a solution that resists change in ph on addition of small amounts of acid or base. It's important to understand that the ph isn't constant, but that the ph changes much less than you would expect it to change if you added the acid or base to a neutral water solution, for example. This importance of buffers is they appear in many biological processes because they are ph sensitive by nature. Biological processes use amino acids and other substances that join together and have acidic ends and basic ends. We all should know the definition of ph. It's the negative log of the hydrogen ion concentration. So in a buffer, the hydrogen ion concentration is maintained more or less. Remember that there is a slight change in ph. The buffer resists change in ph and addition of small amounts of acid or base. Okay, so there are two types of buffer solution, either an acidic buffer or a basic buffer. Acidic buffers are made up of a weak acid and the salt of the weak acid. Good example would be ethanoic acid and sodium ethanoate. A basic buffer comprises a weak base and the salt of the weak base. A good example here would be ammonia and ammonium chloride. Let's deal with acidic buffers first. The acidic buffer is, as I said, we can use it as an example, ethanoic acid and sodium ethanoate. Let's have a look down here. Will see that because the acid, ethanoic acid is a weak acid, it only dissociates very, very slightly into hydrogen ions and ethanoate ions. In fact, we could say as a rough guide, we've got around about 99% of the ethanoic acid molecules and about 1% dissociation of hydrogen ions and ethanoate ions. A In this case, however, the sodium ethanoate, the Salt, it dissociates 100% because it's an ionic salt. And so when we add the sodium ethanoate, we get all of the sodium ethanoate appears as sodium ions and ethanoate ions. Okay, so what's the consequence here? Well, the consequence is that in our buffer solution, our acidic buffer, we have a large amount of the ethanoic ethanoic acid, and we have a large amount of the ethanoate ions. So the buffer mixture contains these two species, ethanoic acid, the weak acid and the iron, the conjugate base of the weak acid in large quantities. And this has, this is important when we add substances. Two, the buffer system. So if we imagine here, this is our. This is a pictorial representation of our mixture. Remember that it's a mixture. So we have a large amount of ethanoic acid represented by this orange bar. And we have a large amount of the ethanoate ions in the mixture. Now of course, the ph is dependent upon the amount of hydrogen ions that are there. Alright, so what we're going to do is this is our buffer mixture, our Britain buffer mixture here we're going to add some acid to the buffer mixture, right, we're going to add a little bit of acid. And let's consider what's going to happen when we add some acid from an external source, some hydrogen ions from an external source. Okay, so first of all then, well, the hydrogen ions in the mixture includes increase temporarily. And what this does is it disturbs the equilibrium at the bottom, the equilibrium between the ethanoic acid molecules and its hydrogen ions and ethanoate ions. And so the ethanoic acid equilibrium moves to the left hand side to restore the value of the equilibrium constant. And it does this by the hydrogen ions interacting with the large reservoir of the conjugate base ions. All right, now let's consider what happens if we add some base. And we've got these two systems here. Remember we've got the sodium ethanoate is completely dissociated into sodium and ethanoate ions. So there's a large reservoir of ethanoate ions and the conjugate, sorry, the acid, the ethanoic acid is only a little bit dissociated. So there's a large reservoir of h. Well, OH ions. When we add them, a base can only interact with the hydrogen ions. Now when they do that, they will remove the hydrogen ions from this equilibrium. They appear in this equilibrium at the bottom, so it gets disturbed. What happens to an equilibrium when it's disturbed? Well, it produces more hydrogen ions by dissociating more of the hair, more of the ethanoic acid. And as we've got a large reservoir of ethanoic acid, we can do that very easily. So we'll add a small amount of base and let's have a look at. Initially, the moles of hydrogen ions temporarily decrease as the hydroxide ions react with them. But then the equilibrium moves to restore the value of the equilibrium constant. And in consequence it replaces the hydrogen ions that have been removed. Now, with the basic buffer, the system works in an identical manner except with different equilibria. This is our example for the weak base, a large reservoir of weak base which is only partially associated in this case with water molecules to produce OH ions. And we've got the salt which is a hundred percent dissociated into ammonium ions. And chloride ions. And so our solution in this case consists of a large amount of ammonia and a large amount of ammonium ions. And this acts as a basic buffer. And we can see how it acts in exactly the same way. The ammonia here is represented by the orange bar. There's a large amount of ammonia in the mixture, the ammonium ions by the green bar, there's a large amount of ammonium ions. And the only equilibrium is this equilibrium at the bottom. The one at the top of course is a full dissociation. It isn't an equilibrium. So the equilibrium at the bottom, if we're going to add some H ions, well, the H ions will immediately remove OH ions. Well, that will disturb this equilibrium here and the equilibrium will shift to the right hand side to make more OH ions. It can do this because there's a large reservoir of ammonia. Let's add the acid. Initially, the hydrogen ions react with the hydroxide ions reduced in their concentration. And then the equilibrium shifts to the right hand side to restore the OH minus ions again. We'll have a look. What happens if we add base? Well, if we add base, clearly the concentration of the OH ions will increase, right concentration of the oh, and if it increases, that will disturb the equilibrium here. Disturbed the equilibrium constant and the way to restore it is by oh minus ions reacting with the large reservoir of NH and the equilibrium of push to the left hand side. Initially, the concentration of H ions increases as base is added. And then the excess ammonium ions from the salt dissociation are able to absorb or react with the new hydroxide ions, making more weak base molecules and water. So this is how buffers, both acidic buffers and basic buffers work. We can do the calculations on this. The calculation just involves using down here. If we have a look, we've got the ethanoic acid dissociation constant. KA is equal to the conjugate base ions multiplied by the hydrogen ions divided by the ethanoic acid. Now, of course, because this is a buffer system, the ethanoate ions and the hydrogen ions are not equal in concentration. Don't get confused with this as just simply a weak acid on its own. The ethanoate ions come from the sodium ethanoate that's been added. Okay, so we just rearrange this equation. Down here we get H plus is KA times ethanoic acid over ethanoate ions and then we substitute our values in. If it's a 0.1 molar solution, well, the ethanoic acid, we assume there's no dissociation whatsoever. That's a Good approximation. And we simply multiply ka by 0.1 and divide by the concentration of the sodium ethanoate, which is. Sorry, there's not sodium. Yes, the sodium ethanoate, because that is the same as the ethanoate ion concentration, which is 0.5. This gives us a value of ka. And from the value of ka, sorry, a value for H, not ka. This is the H equals. This is a value for H. And we then take the negative log to give the ph. It's exactly the same with basic buffers, except in this case we work through kb, our KB equation. KB equation is the conjugate acid, the ammonium ions multiplied by oh divided by the base, the weak base. And again we rearrange this to give oh, etc. You can pause the video, you can have a look through the the calculation. I'm not going to work through the calculation with you. It's exactly the same. Making buffers. Well, we can prepare buffers simply by adding the two components. We can add sodium ethanoate to ethanoic acid. That's one of the ways of making buffers. Again, you can pause the video, you can look at the calculation here to see how to work out how much sodium ethanoate you would have to add to ethanoic acid to get the target ph. And this is the second way. This way is very popular with examiners. A buffer can be prepared by reacting a strong base with excess weak acid. So some of the strong base reacts with the excess acid, making the salt. So we have to do a little bit of a calculation to find out how many moles of salt are produced, how many moles of weak acid are left over. And then once again we plug the values into the KA equation to work out the ph. Once again, pause the video, go through the calculation. Alright, so this is our summary, the summary of buffers. A buffer solution contains excess of either a weak acid or weak base and its salt. The weak acid or base provides a store of molecules that can interact with water and ions that can interact with any components added. The salt dissociates 100%, providing a reservoir of conjugate ions that are then able to absorb added hydrogen ions or hydroxide ions in the case of a basic buffer. And so there's the two systems drawn here. You'll see that similarity. Acidic buffer, large reservoir of weak acid and conjugate base. A basic buffer, large reservoir of weak base and conjugate acid. The only difference of course, is that the PH of the buffers, this will be the acidic buffer, will have a PH in the acidic range. The basic buffer will have a PH in the basic range.