Electrochemistry Electrolysis and Electrolytic Cells Electrolytic Cells Page [1 of 2] In a galvanic cell, we use the potential of chemical reaction to do useful work. We get electricity out, in crude terms. An electrolytic cell is the opposite idea of that. In the electrolytic cell, we’re going to put in energy to drive a chemical reaction in the opposite direction that it would normally want to occur in. So let’s look at a couple applications of electrolytic cells. Now, the simplest type of an electrolytic cell would be one in which the electrolyte – now remember, we need electrolyte in order to conduct charge. The electrolyte, the solvent and the consumables, if you will – the chemicals involved in the reaction – are all the same thing. Well, a convenient way to do that is to take a salt, like sodium chloride, and melt it. Now, remember, ionic materials themselves are not conductive. But when they are melted, they become very conductive, because the ions become mobile. So by taking a molten salt and putting two electrodes in the molten salt and applying a potential sufficient to cause a chemical reaction, we have, essentially, an electrolytic cell. Now, in the case of sodium chloride, what would be going on? Well, at the anode, the only thing we have to oxidize is chloride and chloride would be converted then to chlorine and we get two electrons from that. At the cathode, where the reduction is going to occur, the reduction is going to be sodium plus going to sodium metal. And, again, this is simply because we have nothing else in the pot that we have to worry about. Now this is an important industrial process, in the case of sodium chloride. It is, in fact, how we generate sodium metal and one of the ways that we generate chlorine gas. And, just as an engineering note, calcium chloride often is added to the sodium chloride. And that might seem a little odd. Why throw in an impurity? But the reason is that by adding calcium chloride, we lower the melting point of the sodium chloride. You may recall that if we have a mixture that the melting of the liquid is actually lower than for the pure form. And so this is just another application of that – just a practical application, so that we don’t need quite as high of a temperature in order to melt the material. Now again, this is called a Downs Cell industrial. Now, the next situation that we would encounter is one where we’re in aqueous solution – again, a very convenient place to do electrochemistry. And lets take the simplest situation we can think of. We have to have an electrolyte, so lets use sulfuric acid as our electrolyte and consider possible reactions at the anode. Now, we have two. We have sulfate being oxidized to make . That’s got a very negative oxidation potential here, in this case. And the other possibility we’d have is water going to oxygen. The oxidation of water itself to make oxygen – 4 protons, 4 electrons. And this has got an oxidation potential of minus 1.23 volts. So, which one is going to happen? The answer is whatever is the easiest thing to do. Now remember that the more positive this value is, the more spontaneous that reaction would be. And so, in this particular case, the fact that this is less negative – in other words, that this is algebraically larger than this number – makes this the easier reaction to do. So oxidation of water is going to occur at the anode for this process. Now, let’s point out one thing real quickly. This value’s not actually completely correct, because that’s only true at standard state. And our initial partial pressure of oxygen is not going to be an atmosphere. So that’s going to change this number a little bit, but you’ll notice, there’s such a huge difference in those two numbers that this is going to be the only reaction that we have to worry about. Okay, now what about the cathode? Well, at the cathode, we have really only one reaction to worry about. On the face of it, we could look up two reactions. We could consider the reduction of H+ coming from the sulfuric acid to make . And that, in fact, we know is just the standard hydrogen electrode, essentially. So that’s got a potential of zero. And we have water reacting with electrons to be reduced to make and hydroxide. But if you think about it, this is really describing the same process, because when this occurs – imagine this reaction occurring – the hydroxide that you’d form is immediately going to react with the H+ that we have in solution, making water. And so, overall, we’re essentially doing the same exact chemical process. So our potential is going to just be zero for this. And if you were to look up these values – and you can see over on the side here – that the only difference between these two potentials is that this one incorporates the dissociation of water along with the standard reduction potential for hydrogen. So overall, we’re talking about the exact same process. The only game in town is making hydrogen at the cathode. Just to point out, that the actual value here is going to change a little bit, because we’re not at standard state. But, again, there’s no issue, in this case, because there’s only one reaction we have to worry about. Now, a more interesting question comes when we look at something like sodium chloride. So let’s look at a one molar solution of Electrochemistry Electrolysis and Electrolytic Cells Electrolytic Cells Page [2 of 2] an ionic salt instead of the sulfuric acid electrolyte. And once again, we’ll ask what’s going on at the cathode. Well, at the cathode, now we have worry about the reduction of water to make hydrogen. And we’ve got a potential that describes that. Once again, that potential is not exactly this value, because we’re not at standard state, but fairly close to this value. And then we have to worry about the reduction of sodium – sodium going to sodium metal. Well, this value is so much more negative than this one. Again, algebraically, this number is larger than this one, and so that will be the easiest thing to do. And just to point out here, we could crank our potential by hooking up a bigger battery to the cell. We could get our potential up as high as we want, and we’ll never be able to reduce the sodium as long as water is around, and that’s simply because this is easier to do. And so until this reaction is taken out of play, this will not happen no matter what the potential is on that battery. Okay, so what’s going on at the anode? Well now, this is more interesting question. At the anode, we have the oxidation of water, to make and we have an oxidation potential that describes that. We also have the possibility of the oxidation of chloride to make chlorine. And we can look up a reduction potential, reverse it to make an oxidation potential, and compare those numbers. Now we have an interesting situation. These values are very close together, and that, in consideration of I’m not at standard state, where these numbers going to be able to change a little bit from those listed values, I’m going to have an interesting situation where, at very high concentrations of chloride, this becomes the dominant reaction, as governed by the Nernst equation. But at low concentrations of chloride, this reaction will take over. And we’ll see oxidation of water. So as long as we’re careful to keep our chloride concentrations high, then this becomes the dominant reaction and that’s useful industrially. We look at the overall reaction, combining that anode reaction with the previous cathode reaction. We end up with this chemical reaction, conversion of chloride and water to make and and hydroxide – this is now sodium hydroxide, if we use sodium chloride to start with. And so, this is a process called the core alkali process, and it’s one of the major sources that we get chlorine gas and sodium hydroxide from. So, again, electrochemistry buys us a lot, as long as we’re aware of competing half-reactions and can use the Nernst equation, or principles of the Nernst equation, to select between them. Finally, let’s talk about the possibility of, instead of generating an element like chlorine, generating an element that’s metallic, and using that to coat a material. Imagine taking a block of iron, for instance, and dipping it into a solution of electrolyte, applying a battery to it, and actually coating out a different metal, perhaps to protect the iron. We talked about corrosion. Well, we can accomplish that, provided that the reduction to make the metal from the metal salt happens easier than does the reduction of water. So, once again, we consider competing reactions at the cathode. And, just while I’m thinking of it, just let me hold this up as an example here of something that’s chrome-coated. This is an example of something that’s silver-coated. We know about bumpers. We know about gold-coated jewelry. Lots and lots of things in our lives are metal coated.