Chemistry: Electrochemical Cells

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When last we spoke, we had a brilliant idea. We were going to take advantage of the fact that copper reacted with silver plus to make copper 2 plus and silver and we were going to harness the electrical energy involved in this reaction. The idea was if we could separate the copper and the silver from each other, we'd be able to pass the electrons through a wire rather than directly through a solution and be able to use that electricity to run a motor or light up a light bulb or do something. So, we wanted to keep things simple so that we could keep things at standard state. We went ahead and designed something where we'd have 1 mole per liter of solution of the copper 2 plus and the silver plus, as well as chunks of copper and silver. So everything was at their standard states. We'd hook up a wire between the copper and silver and the sad news was that this didn't work. And the question I left you with was why not?
Well, I know you haven't got much sleep since last time we talked, but the reason is that as electrons start to move from this cell to this cell, and copper starts to become copper 2 plus, and the first few atoms of silver start to precipitate out to make silver metal, what happens is that the accompanying ions start to build up charge. In other words, this wasn't just silver plus; it was silver plus with a nitrate. Over here, the copper, once it formed copper 2 plus, that copper 2 plus doesn't have an anion accompanying it. What we need to do is somehow get the nitrate anions, the negatively charged anions, over to this cell to account for the charge that we're building up here. And that's the reason why our idea, as we originally devised it, would never work. We could pass the first couple of electrons and then that's it. We're going to build up charge, and remember, one of the fundamental things we know about nature is nature does not like separation of charge, and that's exactly what we're doing.
So, we need to prepare a path to get nitrate from this cell back over to this cell. But of course, we also want to not mix this copper solution over here. Otherwise, we're going to just get the chemical reaction that we saw in the beaker a few moments ago. So, the answer is - kind of the practical solution to this is - we want to have some type of a permeable membrane, but one that slows down the fusion enough so that we don't let the copper get over here. Now another way to accomplish the same thing is if we have a path in which ions can flow through but the point would be that the path would be long enough that we don't allow the copper 2 plus that's formed to get over into this other side, because then the copper and the silver would react directly.
So what's done in practice, then, is we use something called a salt bridge. Imagine a piece of glass tubing used here to connect the silver and the copper cells, or half-cells. And then at the ends of the tube, we'll put in a plug of cotton or something to prevent solution from easily flowing in and out. But, again, the cotton has small channels where ions can flow. Finally, what we're going to do is put in an electrolyte in this solution so that we allow for, if not copper 2 plus ions, at least the sodium cations that are in this electrolyte - like it might be sodium nitrate, for instance - to start to migrate towards the negative charges here and the nitrates to start to move into the salt bridge while other nitrates move out of the salt bridge.
So the details of this are not important. What is important is it allows us to move charges from this cell to this cell in the sense that we balance out the charges that were moving from this electrode to this electrode. In doing that, we complete a circuit and, indeed, we can get work from this circuit. And, in fact, this is essentially the essence of a battery. This is a 9-volt battery here. We'll talk more about batteries - as an example - something called a dry cell. It's not a wet cell like I'm showing you here. We'll talk lots more about what different batteries there are and how they're designed a little bit later. But the basic idea of what we show on paper here is exactly what a battery is. It's harnessing chemical energy to use in the form of electrical work, in this case.
So let's define our terms here again. This is called a galvanic cell. And a galvanic cell is a cell that will run spontaneously. You'll see we'll contrast this in a moment. It's a cell that can run all by itself once we hook it up. That just simply means the chemical reaction - the net chemical reaction we're doing - is a spontaneous reaction. We know all about that now. And so, we will define that anode as the electrode where oxidation occurs. So in this case, remember, the oxidation is copper metal going to copper 2 plus, and so that is where the anode is. In this case, that's what we defined. And normally, by convention, people will have the anode on the left side, although, I'm not a big subscriber to that convention. I think it's kind of pointless to remember something like that. Just remember that anode always is where oxidation is. That's the important thing to remember.
Now, in contrast, what's going on over here is a reduction. We call that the cathode. So the definition of a cathode is where the reduction's occurring. That's the electrode where the reduction's occurring. And, let's see, what else can I tell you about the cell? Pretty much, that's it. We can either use this electrical current to do work or, what we'll do a little bit later is take out this motor - take away the load - and just simply measure what the potential is to do the work. And we can actually measure a voltage across this, so we'll get to that in just a moment. For now, I just want to introduce you to the galvanic cell, and we'll compare this to an electrolytic cell.
An electrolytic cell is the opposite idea. An electrolytic cell is a cell where you're running the current in the opposite direction that it's normally inclined to go. So if we take our copper and silver cell, just like we had before, but we replace the motor with a battery such with enough potential to force electrons to go in this direction rather than in the opposite direction. We still need our salt bridge to make sure that we complete a circuit, but then let's look what's going on. If we're forcing electrons out, what we're doing is pulling electrons now out of this electrode. That means that now the oxidation occurs over here. So this becomes the anode now. We have silver metal going to silver plus plus an electron. The electron is being pulled out into the battery here. And notice that the anode is now on the right side, in an electrolytic cell. And again, all I care about, at least, is that you know that an anode is where oxidation occurs. That is still true, regardless of what kind of cell we have.
Over here is the cathode now. This is where the reduction happens. And what is the reduction in this case? Copper 2 plus plus 2 electrons goes to copper metal. So we're reversing the process and we're plating copper metal out over here and we're chewing up the silver in the process. Now, who in their right mind would want to do this and consume silver in order to get copper metal? Probably nobody. You could think of a lot of other materials that you could use instead of silver as a way of plating copper out. But the basic idea here is actually used industrially. This is the essence of electroplating, and we'll have more to say about that. When people talk about spoons that are coated with silver or gold-plated jewelry and so forth, it's this basic idea of how that is done.
So, again, we'll come back to that later. Right now, it's just the definition of the terms. Electrolytic cell, again, is something running the opposite direction that it normally wants to run in, whereas the galvanic cell, it was running the direction that would be the spontaneous reaction. So we're going, seemingly, against nature here - apparently, in violation of Gibbs free energy, but actually not, because we're forcing it, now, with an even greater free energy in the opposite direction, if you will, with this battery forcing this situation to go in reverse.
So, again, that's an electrolytic cell, a galvanic cell, and indeed, our basic notion that we can take certain chemical reactions, split them into pieces - what we'll refer to as half-cells - and get useful work out of them. Our next step, then, is to actually remove that motor and measure the potential to do the work rather than actually do the work. That, as you might guess, is going to be related to the driving force of the reaction. We're going to make a fundamental tie between Gibbs free energy, which describes how reactions desire to go to equilibrium, and the potential that we can measure across those wires before we connect it up and have the reaction occur.
Electrochemistry
Galvanic Cells
Electrochemical Cells Page [2 of 2]