Lecture 18. Electrochemistry Pt. 3.
Arasasingham, Ramesh D.
 UCI Chem 1C is the third and final quarter of General Chemistry series and covers the following topics: equilibria, aqueous acid-base equilibria, solubility equilibria, oxidation reduction reactions, electrochemistry; kinetics; special topics. 
 Okay, so were there any questions about the material from last week? Yes, go ahead. The lecture from Friday wasn't posted yet. Okay, it should be up, so let me check and let you guys know. So the lecture from Friday, I think, is up on the website, isn't it, Sean? Okay, it is up. Can you check and let me know? Send me an email and I'll forward it to him. Okay? Anything else? Yes? If you're an acid, yes. If you're balancing redox reactions and acid acidic solution, it's a lot easier to put H plus rather than H3O plus. Alright? But technically speaking, it's more accurate to put H3O plus, okay? But it's a lot easier to balance it if you just have H plus. Got it? Okay, so today we're going to move on to the actual topic which deals with electrochemistry. And we said electrochemistry deals with redox reactions. And so let's take an example of a redox reaction, the one written here where you have zinc metal reacting with copper sulfate to give you zinc sulfate and copper. Now this is a spontaneous reaction. Remember last quarter when we looked at thermodynamics, when a reaction is spontaneous, it means that it has a natural tendency to occur on its own without any external influence. And the thermodynamic quantity that we use to describe a spontaneous reaction is free energy. But if the change in free energy for that reaction is negative, we know that it's spontaneous. So we're looking at a spontaneous reaction and this is written in the form of the general reaction or the overall reaction. Now in many textbooks they use the term molecular equation, but as all of you can see, we don't have molecules involved here. We have ionic compounds, alright? Therefore I'm just going to call it the general form of the reaction, okay? Now if we wanted to write the complete ionic equation, then we would write out all the dissociated species. So if you start with zinc, that's metallic zinc, so that does not dissociate in water. But we know copper sulfate is a soluble salt. How do you know? Solubility rules, okay? The solubility rules tells you that copper sulfate is soluble and if it is a soluble salt, we know that it dissociates completely to give me copper 2 plus aqueous plus SO4 2 minus aqueous. So in reality this is what the reactants look like in solution. Now on the product side, now we have zinc sulfate which is a soluble salt. So the zinc 2 plus and SO4 2 minus would be dissociated and then you end up with copper metal. So that would be the complete ionic equation. And if you look at the complete ionic equation, you can see there's sulfate at the beginning, there's sulfate at the end. So in this reaction, the sulfate anions are the spectator ions. And now if we cancel out the spectator ions on either side, then we end up with the net ionic equation. And that would be zinc solid plus copper 2 plus aqueous, giving me zinc 2 plus aqueous plus copper solid. So now you can see that this, the net ionic equation actually describes the net event that takes place. So this tells you exactly what's going on. So the net ionic equation kind of summarizes exactly what's going on. And if you look at this, one way to figure that out is to assign oxidation numbers. So this is a metal in its elemental state, that would be zero. This has a plus 2 oxidation state. This has a plus 2 oxidation state and this is copper metal. So you can see that if I take zinc, the oxidation number of zinc goes from zero to plus 2. And we know that when the oxidation number goes up, what is that process called? Oxidation. And if you look at the relative change, it changes by 2. And since oxidation represents loss of electrons, we know that two electrons are lost because the oxidation number changes by 2. All right? Now let's look at copper. If you look at copper, it goes from plus 2 to zero. So copper, the oxidation number decreases. And if it decreases, we know it's a reduction. And it changes by a magnitude of 2. And therefore, we know two electrons are gained. So two electrons are lost and two electrons are gained. And therefore, you know that two electrons are being transferred in this reaction. And if you wanted to write this as half reactions, remember we said half reactions don't really take place. It's just a convenient way for us to keep track of what's going on. And a half reaction is purely a mental process for us to kind of figure out what's going on. Okay? So if you want to write half reactions, then the oxidation half reaction would be where zinc solid goes to zinc 2 plus aqueous plus two electrons. And this is the oxidation half reaction. All right? It's oxidation because the electrons on the product side, electrons are being lost. Our electrons are being produced here. So that will be the oxidation. And then on the other side, we have copper 2 plus aqueous picking up the two electrons that are lost to give you copper solid. And this would be the reduction half reaction. And in this case, the electrons are on the reactant side. That means they're being picked up in the reaction. So electrons are being gained in the process. And therefore, that is the reduction. Okay? Now just like in acid-base reaction, remember acid-base reactions involve proton transfer. Redox reactions involve electron transfer. Just like a proton is lost and a proton is gained in an acid-base reaction, in a redox reaction electrons are lost and electrons are gained. So just like in acid-base reactions, remember we talked about conjugate acid-base pairs and we always write down conjugate acid-base pairs. Likewise, we can write redox couples. So just like conjugate acid-base pairs, I used to describe acid-base reactions. Here we use the term redox couple. And in the redox couple, you always write the oxidized form first, then a diagonal line, and then the reduced form. So always you write it in that order. So in this reaction, you can see there are two redox couples. All right? One is zinc and the other is copper. You always write the oxidized form first, so that would be zinc 2 plus, which has oxidation number of 2, and then zinc solid. The physical state would be aqueous. And then on the other side, we have copper 2 plus aqueous and copper solid. So in this instance, you would always write the redox couple. And you can see if you have an oxidation reduction reaction, you will always have two pairs of the redox couples. All right? So the redox couples here are always written in the form of the oxidized form, a diagonal line, and then the reduced form. Is that clear? So if any reaction you should be able to identify the redox couples, just like we did for conjugate acid-base pairs. Now in the case of conjugate acid-base pairs, it's easy because the acid form has an extra proton. The basic form has lost that proton. So you can just look at that and see the proton lost, and you can identify which is the acid and which is the conjugate base. Here it's not directly apparent. You have to look at oxidation numbers to figure out which is the oxidized form and then which is the reduced form. All right? And write it in that form. Now the reason we're looking at all of this is ultimately in electrochemistry, what we're interested in looking at are called electrochemical cells. So we're going to start looking at electrochemical cells. And let's define what they are. So ultimately we're going to look at what we call galvanic cells, or sometimes the term voltaic cells are used. And both the galvanic cells or the voltaic cell is called electrochemical cell. So let me define what an electrochemical cell is. An electrochemical cell is a device in which an electric current, all right? Now all of you know electric current is what? A flow of electrons. A flow of electrons produced by a spontaneous chemical reaction or is used to bring about a non-spontaneous reaction. I know what that buzzing sound is. I know it's annoying, but it's a whole annoying. So let's go back. So we're looking at galvanic cells. And galvanic cells, voltaic cells are called electrochemical cells. Overall electrochemical cells are devices that we make up, we build. And they can be a few types. One is that you take advantage of a spontaneous reaction to generate a flow of current, all right? Or you can take an external source of current. So you plug it into a plug, for example, that's an external source of current and you use that to drive a non-spontaneous reaction. So essentially we're interested in looking at spontaneous reactions and non-spontaneous reactions, all right? We're going to start by looking at those that take advantage of a spontaneous reaction. And one that takes advantage of a spontaneous reaction is called a galvanic cell. Okay, so a galvanic cell, for voltaic cells, is an electrochemical cell in which a spontaneous chemical reaction is used to generate an electric current. All right, so what we're going to do is we're going to start by looking at spontaneous reactions. And so if you take a galvanic cell or a voltaic cell, this is an electrochemical device. It's a device that you construct and it takes advantage of a spontaneous chemical reaction to generate electric current. So bottom line, what are galvanic cells or voltaic cells? So these are batteries. So in electrochemistry we call them galvanic cells or voltaic cells and the simple idea behind this is that if you want to construct a battery and you know that there are a whole host of batteries out there, they're lithium batteries, there are different kinds of batteries that are used for different purposes. Now all of these are galvanic cells. The underlying principle is the same and they take advantage of a spontaneous reaction. Okay? So let's take an example. So we just looked at the fact that if you take this reaction where you have zinc solid and copper 2 plus aqueous, we know that goes to zinc 2 plus aqueous plus copper solid and this is spontaneous. All right? And that means delta G is negative. So we know that if I take a chunk of zinc metal, all right, and take a solution of copper sulphate, copper sulphate is a soluble salt. So you take a spoonful of copper sulphate, put it in water and stir it up and then put a chunk of zinc metal inside. What's going to happen is this is a spontaneous reaction. There's a natural tendency for this reaction to take place and you're going to have the zinc metal dissolving as it reacts and then at the bottom you're going to see copper metal depositing at the bottom, all right, as this reaction progresses. And at the end it's going to establish equilibrium, all right? Every reaction, remember, ultimately goes to equilibrium. So there comes a point in which the reaction will start slowing down, not really slowing down, but the amounts of the reactants and products start decreasing and there comes a point where the reactants start decreasing, the products are going up and there comes a point when the amounts of both stop changing and at that point we know the equilibrium is established, okay? Now when we analyze this reaction we know two electrons are being transferred, all right? So that means if I have one atom of zinc and one copper ion, we know two electrons will be transferred, all right? So if you just have one atom and one ion of each, you're going to have two electrons transferred. Now if I have billions upon billions upon billions of zinc atoms and billions upon billions upon billions of copper atoms, now how many electrons will be transferred? Twice as much billions, so twice as much billions upon billions of electrons will be transferred. Now if I take one mole of zinc and I have one mole of copper, we know two moles of electrons will be transferred. You see that? If you have amounts of both, you're going to have a tremendous amount of electrons being transferred, okay? Now the idea behind batteries is quite clever. The idea is that if you take just the two of them and the intimate contact with each other, the electrons are going to be transferred and the system will go to equilibrium. Now what if we could separate the systems so that now we're going to take advantage of this spontaneous reaction, but if we can physically separate this from this, all right? And then if we connect them through a wire, can we force these billions upon billions of electrons, can we force it through the wire? And in that way, you can now generate a current, all right? So the idea is instead of having the two reactants in intimate contact with each other, where the electrons will be transferred from one to the other, rather we want to construct a device where we're going to take advantage of the fact that there is a spontaneous reaction, that electrons are going to go from one to the other, but now we're going to force these electrons through a wire, all right? Physically separate them and force it through a wire so that now you have electron flow through the wire. Got it? And remember, electron flow is what? You can get a current, okay? So that means we have to construct a device where you have all the elements for this process to take place. So we have to separate the zinc from the copper. In order to separate that, what we deal with is where we have the zinc metal, and we're going to call the zinc metal, we're going to call this the electrode. And this zinc metal is going to be dipped in a solution of zinc sulfate, okay? Usually this is about one molar concentration, and so since this is soluble, what you'll end up with is zinc 2 plus equus and sulfate 2 minus equus, each of which will be one molar. So on one side, you have the reactant zinc, and so that's one half of that redox reaction, okay? On the other side, you're going to have copper, so this will be copper metal, and over here I'm going to have zinc metal. This drawing is really bad, but I'll show you a nice picture later, okay? And this is dipped in copper sulfate. So this is the copper electrode, and we have copper sulfate equus, AQ, and this will be one molar typically, and this dissociates to give you copper equus and sulfate because it's a soluble salt, okay? Now I'm going to link these through a wire, and if you wanted to just to make sure that there's a current flowing through, we can link it through a voltmeter. So this is a voltmeter, alright? And so now I've physically separated the oxidation half reaction from the reduction half reaction so that we have the two reactants separated, and we're going to force the electrons to that wire, but if I'm going to force the electrons to flow to that wire, all of you recognize that I have to complete the circuit, okay? The circuits are completed. In order to complete the circuit, what we have is what we call a salt bridge. So this is a salt bridge, and a salt bridge essentially is a gel that's in that U-shaped glass tube, and over here I'm going to put a cotton plug to prevent the gel from seeping out, alright? And in here, this gel has KCl equus. In other words, it has potassium ions equus plus Cl minus equus, alright? And I'm now going to switch and show you a pretty picture because it's the same thing that I drew, alright? So now I'm going to switch to a picture that kind of shows you. So let's start with this picture in the middle. So we have the two electrodes. So we have the zinc electrode and zinc metal dipped in zinc sulfate, and that dissociates to give you zinc ions and sulfate ions. On the other side, we have separated the other reactant, which is copper, and that is copper metal dipped in copper sulfate, and that has copper ions and sulfate ions, and the two are connected through a wire, and we're going to attach it to a voltmeter so that we can measure to see where there's a flow of current or not, okay? Now in order to complete the circuit, what we have is a salt bridge. So here is the salt bridge, and the salt bridge is essentially a gel that's placed in there that has ions in it, dissociated ions, and so you have potassium ions and chloride ions dissociated in that, and then at the bottom of the glass YouTube that has the gel, you have cotton plugs on either side just so that your gel doesn't come flopping out, all right? Now when we complete the circuit, we know that when you have flow of electrons, you have a current flowing. Another, a different way in which you can have current flowing is also to have motion of ions, all right? If you have movement of ions too, you can have a flow of current, all right? And so what's happening here is that you have a flow of electrons that go from this to this, but now you complete the circuit because now you can have ions flowing from this to this, all right? And so overall what we've done is we've completed the circuit. So there's a complete circuit so that now you can have electrons flowing from one side to the other, got it? So this is simply put, this is a very rudimentary battery. This is the idea behind batteries, all right? Now if you look at this, at the left hand side is where the oxidation takes place. So in this, on the left hand side, this is where, we call this the anode, and the anode is where the oxidation occurs, okay? And the sign for the anode, the anode has a negative sign, all right? And what happens at the anode is electrons are pushed into the wire, all right? So if you look at the left hand side, usually it's convenient to draw these out so that you put it in alphabetical order. One side is anode, the other side is called a cathode, and therefore it always comes first, all right? So at the anode, you always have oxidation taking place, and when oxidation occurs, what happens? You have electrons being produced, it's electron loss. So when the oxidation takes place, you have lots of electrons being produced, or lots of electrons being lost. Remember those electrons are going to be channeled through the wire. So at the anode, electrons are being pushed into the wire so that you have a flow of electrons, okay? Now on the other side, you have the cathode. So and this is where reduction occurs, and this has, is symbolized by a positive sign, and we say electrons are pulled in from the wire, all right? So the wire is connected to the copper electrode, and so the cathode is where electrons are being pulled in. So any electrochemical or galvanic cell has two components. If you take a galvanic cell, it's got an anode, because we've separated the two reactions. If they were in contact with each other, now the electrons will be transferred one to the other, if they're in contact with each other, then that's not useful to us. So we're taking advantage of a spontaneous reaction that we know has a natural tendency to occur without any external influence. We're going to physically separate the oxidation half reaction from the reduction half reaction, and we're going to connect the two through a wire so that the spontaneous reaction still occurs. So now you're forcing those electrons, those billions upon billions upon billions of electrons, now are being channeled through that wire from one to the other, all right? So wherever you have the oxidation, we call that the anode. The anode always has a negative sign, and then you know that at the anode, electrons are being produced, and those electrons are being forced through the wire. They're going to push through the wire. And at the other end, the flow of electrons comes into the cathode. So the cathode is where the reduction takes place, where electrons are being gained, and so the wire, the electrons are taken up, and that's where the reduction takes place, and so you have a pull of electrons. So simple put, what you have here is a push and pull of electrons. You have one terminal, you're having electrons being pushed to the wire, and at the other terminal, you're having electrons being pulled in, okay? And I kind of like this picture because it summarizes everything that's going on. So if you look at this, and if you look at the bottom here, you can see that this is where the oxidation takes place, and zinc is oxidized to zinc 2 plus. So the net reaction is zinc is going to zinc 2 plus plus 2 electrons. At the oxidation point, two electrons are being produced. If you have the ability to magnify what's going on, this is the zinc electrode, which is a metallic solid, all right? So the zinc electrode will give off two electrons, and it becomes zinc 2 ions, and when it becomes an ion, it dissolves in solution. You guys see that? So what's happening at the anode is that the metal is actually dissolving because the metal is giving up, the metallic solid is giving up two electrons and pushing it up through the wire and dissolving as zinc 2 plus. So at the anode, what's going to happen is if you start with a solid metal strip of zinc and you weigh it before you start, you connect everything up and complete the circuit, if you weighed before and then you weighed afterwards, what's going to happen to the mass of that strip, metal strip? It's going to go down. Can everybody see that? Because at the anode, oxidation takes place, and that means the metallic solid is now converting to zinc 2 plus and the zinc 2 plus dissolves in solution. All right? Now if you go to the other side where the reduction takes place, now here's the reduction half reaction. So the electrons that come into the copper electrode are being pulled in, and the two electrons react with copper 2 plus, the copper 2 plus ions in solution, and you form a copper solid. So what happens at the cathode is you're having more and more copper depositing on that electrode. So if you weigh that electrode before anything happens, and if you weigh it afterwards, you'll see that its mass goes up. All right? And so what's going to happen is at the anode, your electrode will start dissolving and getting smaller and smaller and smaller, whereas here where you have the dissolved copper ions picking up the two electrons. So they're going to pick up the two electrons to give you copper solid, and these will deposit on the electrode, and therefore the mass of the electrode increases. All right? And this will continue until the reaction reaches equilibrium. And when the reaction reaches equilibrium, now you don't have any net flow of electrons. All right? And that's what we mean when we, in real life, we know that all batteries go dead after a while. And when a battery goes dead, what does that mean? That equilibrium has been established. All right? Once equilibrium is established, now delta G equals zero. All right? And therefore there's no flow of electrons, and so the battery goes dead. How do we know that there is an electron flow? Use a voltmeter to measure it. Now, have you guys done a rudimentary electrochemical cell in the lab yet? This week. Okay. So now, one of the things in the lab is if you attach, you set up your electrochemical cell, you know what your anode is where the oxidation is. You set up your cathode, you know where the reduction takes place, and you hook it up to your voltmeter, and your voltmeter reads a negative sign. That means that you reverse the terminal, because when you measure the voltage, the voltage should always be positive. All right? And so when you hook it up to the voltmeter and complete the circuit, and if you see a negative sign on your voltmeter, that means you reverse the terminals. That means you, you know, remember the anode is the negative terminal, the cathode is the positive terminal, and therefore you have to hook it up appropriately. Got it? Now we said lastly that you have to complete the circuit, and we're going to use the movement of cations and anions to complete the circuit. All right? And if you look at this, what happens is that not only are you completing the circuit and this movement of cations and anions help the flow of current through the system, but also they also help maintain charge balance. So now if you look at the zinc side, remember zinc solid is going to zinc 2 plus. So you're starting with something that's electrically neutral, and in that solution you're having more and more cations. Remember everything in nature has to be electrically balanced. So if you're forming more and more cations, you need anions to counterbalance that charge. And so what happens is in this gel, to counterbalance the extra zinc ions that are coming in, the anions start flowing into that solution to balance the charge. So you have chloride ions moving in, migrating through that gel, they'll migrate and balance the charge as well. So not only do the flow of the ions generate a current, and you can complete that circuit, but also they help maintain charge balance. Now on the other side, can you see that you have copper ions depositing? So you have copper 2 plus cations going to something that's electrically neutral. So in this solution, if all of your cations become neutral, you can see you have lots of anions in there that don't have a counter charge to balance that charge. So to balance, because in that solution you have lots of SO4 2 minus, the sulfate anions are there, only the cation becomes electrically neutral, and therefore you need cations to counterbalance those sulfate ions. And so to balance the charge, you have migration of potassium ions that migrate through the gel into that. Got it? And so that describes a galvanic cell. Yes, go ahead. The sulfate is a specter ion, okay? But you need sulfates because everything in nature has to be electrically neutral. So if you have a cation, a cation always has to have an anion that balances the charge. Do you see that? To begin with, before any reaction takes place, you have the same number of zinc ions and the same number of sulfate ions in here, and you have the same number in this solution, the initial amount of copper ions will equal the sulfate ions. Got it? But once the reaction starts proceeding, you know that your zinc is, your zinc metal becomes zinc 2 plus, so in this solution, you're going to have more and more cations building up. So to balance the charge of the increasing number of cations, you need anions to migrate in there to balance the charge. Is that again? You have to have the sulfate. In nature, you can't have cations without anions. Do you understand that? So take sodium chloride. You can't have, if you take sodium chloride, which is an ionic salt, you cannot have sodium cations alone. Did you understand that? So you need chlorides to balance that. Yes? Is that again? You can't have a salt bridge. Salt bridge. Salt bridge. Without the salt bridge. Is that what your question was? Okay. Without the salt bridge, the circuit won't be complete. Do you understand that? You need the salt bridge because for electrical flow, your electrical circuit has to be completed. All right? And so if you didn't have the salt bridge, nothing would happen because you have to have the circuit completed. Okay? Sorry. I miss, I heard the sulfate, not salt bridge. Okay. So we've looked at an example and this particular cell is called the galvanic cell. The galvanic cell has a particular name and this is called the danial cell. So if, so we call this the danial cell, so let me just, so if your reactants are zinc, metal and you have zinc, zinc sulfate and copper, copper sulfate, this combination of a galvanic cell has a specific name and this is called the danial cell. All right? Now you can have different types of electrodes. So let's take another example. So if you want to describe another example of a cell, let's take another electrode. So we looked at electrodes and we've looked at zinc electrode. We've looked at copper electrodes. All right? So in the danial cell, the zinc electrode forms the anode, the copper electrode forms the cathode and you have a flow of electrons. Okay? Now you can also have another electrode called the hydrogen electrode. So let's take a look at what it means when we say we have a hydrogen electrode. And let's say your hydrogen electrode is the cathode. So if your hydrogen electrode is the cathode, remember that the cathode, does oxidation take place or reduction take place? Reduction. All right? So the reduction reaction that you see at the cathode will be 2H plus equus plus 2 electrons giving you hydrogen gas. All right? And so if you wanted to set this up, your H plus comes in the solution in the form of let's say HCl equus and usually the concentration is one molar. So you can see in this, it's a strong acid, it dissociates completely. So in that solution you have H plus ions. All right? Now the product is hydrogen. So we need something. Remember we need a wire to bring in the electrons into, remember where the reduction takes place, electrons are being pulled in. So you need a flow of electrons coming into that solution. And so what we do is we have an inert metal like that and this whole, the wire as well as this is platinum and this is an inert metal. So the reason it's inert is it doesn't react. All it does is just serve to deliver the electrons into the solution and you have a test tube that looks something like this. So you have a test tube which is open at one end like that and you have hydrogen gas as well. And so essentially what happens is the hydrogen gas comes through here and then it'll bubble out on these sides. You'll see bubbles coming out and this would be the cathode. All right? So if you wanted to look at a setup where your cathode was now a hydrogen electrode, this is what it would look like. So here on this side we have the anode and the anode is the zinc electrode. The cathode is a hydrogen electrode. So essentially what you do is you have hydrogen gas at one atmosphere, you have one molar HCl, your platinum, the platinum metal is just used because it's an inert metal. You need something to deliver the flow of electrons into that solution so that this reduction occurs and so an inert material like platinum is used. All right? And you can see the bubbles of hydrogen gas. So if you wanted to describe what's going on in this cell then you would say that at the cathode, so at the anode you would have the oxidation taking place which is the zinc solid going to zinc 2 plus equus plus 2 electrons. At the cathode you would have the 2H plus ions picking up 2 electrons to give you H2 gas. And so your net ionic equation would be where you have zinc solid plus 2H plus equus giving you zinc 2 plus equus plus H2 gas. Did you guys get that? So overall the net ionic equation that takes place in this galvanic cell is shown here and you've identified what the anode is, you've identified what the cathode is and then how many electrons are being transferred? There are 2 electrons being transferred in that setup. Now you could also have the other way around. So now let's say let's take a different cell and in the other cell let's say the oxidation. So what we have is the anode is where you have the oxidation. So now let's say we have the hydrogen electrode as the anode. So the hydrogen electrode is the anode, you're going to have hydrogen giving you 2H plus equus plus 2 electrons. This would be the oxidation. At the cathode you would have copper 2 plus equus pick up the 2 electrons to give you copper solid. So this is the reduction. And now we've looked at an example where the hydrogen electrode is the anode and now the copper electrode is the cathode. So we took the Daniel cell and we took one half of it and replaced one half with the hydrogen electrode on either side. And you can see that the net reaction would be, so the net ionic equation would be where you have hydrogen gas plus copper 2 plus equus giving you 2H plus equus plus copper solid. So the overall reaction that takes place in the second example where now your hydrogen electrode is the anode where the oxidation takes place and now copper is the cathode then the overall reaction would look like that. So that's what this picture shows you. So the two examples that we looked at here is remember usually when you draw these out you're customary to always put the anode first and then the cathode second. Here you have the zinc electrode here and over here you have the hydrogen electrode. So this is the cathode and you have a flow of electrons and the circuit is completed with the salt bridge. Now over here this is the other. Now you have a copper electrode which is the cathode. Your hydrogen electrode is now the anode so now this is where the oxidation occurs. This is where electrons are being pushed out and it goes around and the electrons on this side are being pulled in. You see that? These are all examples of galvanic cells. A galvanic cell is an electrochemical device that takes advantage of a spontaneous reaction. It has to be a spontaneous reaction if you're going to have a flow of electrons to the wire. And so galvanic cell is another name for batteries. This is the underlying principle for all batteries. So now we're going to look at the last few minutes. We're going to look at what we call cell notations. So this is chemist love shorthand notations. And so we have developed this system of notations that we use to describe cells, particularly galvanic cells. So we're going to look at cell notations. And let's start with the Daniel cell. We know in the Daniel cell, the reaction that we're looking at, the net reaction, is that you have zinc solid plus copper 2 plus ions giving you zinc 2 plus equus plus copper solid. So that isn't it. And so in this reaction, we looked at the fact that the redox couples are, we have the oxidized form first, so the zinc 2 plus equus, we use a diagonal line and then zinc solid. And we have copper 2 plus equus, copper solid. That's the redox couple. All right? Now if we want to describe the cathode for this reaction, let's start with the anode because always in alphabetical order. So let's start with anode. Now when you write the anode, and this is called the notation for the electrode. So this is, we're looking at the electrode. You always write the reactant first and then the product. All right? And if there's a vertical line like that, now it's different from a diagonal line. It's very, very important that you draw vertical lines and diagonal lines and don't mix up the two. All right? If you draw a vertical line like that, it means that there, the reactants and products are in two different phases. So if you have something that's dissolved in aqueous solution and a solid, or if you have aqueous and a gas, or you have a gas and a liquid, all right? These are two different phases and what we say is that there's an interface between them. All right? So if you have a vertical line, there are two different phases. Or if you always write the reactant first and then you put a comma and then product, this means same phase. All right? So if I wanted to write the Daniel cell, my anode would be, my reactant for that is zinc. I'm going to put a vertical line because we're going from zinc to zinc 2 plus aqueous. All right? So that describes the anode. Now if I want to describe the cathode, I would write the reactant. Vertical line because again, you're going from an aqueous solution to a solid so there's a change in phase would be that. All right? Now we describe the cell, so the cell notation is always anode first, cathode last. All right? So just remember it as AC. All right? It's an alphabetical order. Just remember that. Anode always comes first. All right? And therefore, I would start with the anode. Remember with the redox couple, it's always the oxidized form, diagonal sign, reduced form. When you draw the symbols for the electrodes, the notation, it's always reactant first, product. And if there's a phase change, you put a vertical line between them. Okay? So now if I combine them to give the cell notation, what you do is you start with anode and you know the anode is that, so vertical line, zinc 2 plus aqueous. Now I put a double line and this double line refers to the salt bridge which completes the circuit. And then on the other side, we have the cathode reactant first vertical line product. So what does this symbol convey? To chemists, this symbol represents this picture. So every time you see this, that's what this picture represents. So the shorthand notation is a very convenient way to convey to everybody. That symbol, whenever you see those sets of symbol, you have to have a mental picture of what your galvanic cell looks like. And so that translates to this. So you know that the anode has a zinc electrode dipped in zinc 2 plus, alright? And that it has a salt bridge and on the other side, you have a copper dipped in copper 2 plus and it signifies this picture. Got it? Now at the beginning of next class, I will take a few more examples of writing the cell notation and show you how to write cell notations to describe what the galvanic cell looks like. Okay? So I will stop there for today.
 Okay, so were there any questions about the material from last week?  Yes, go ahead.  The lecture from Friday wasn't posted yet.  Okay, it should be up, so let me check and let you guys know.  So the lecture from Friday, I think, is up on the website, isn't it, Sean?  Okay, it is up.  Can you check and let me know?  Send me an email and I'll forward it to him.  Okay?  Anything else?  Yes?  If you're an acid, yes.  If you're balancing redox reactions and acid acidic solution, it's a lot easier to put  H plus rather than H3O plus.  Alright?  But technically speaking, it's more accurate to put H3O plus, okay?  But it's a lot easier to balance it if you just have H plus.  Got it?  Okay, so today we're going to move on to the actual topic which deals with electrochemistry.  And we said electrochemistry deals with redox reactions.  And so let's take an example of a redox reaction, the one written here where you have zinc metal  reacting with copper sulfate to give you zinc sulfate and copper.  Now this is a spontaneous reaction.  Remember last quarter when we looked at thermodynamics, when a reaction is spontaneous, it means that  it has a natural tendency to occur on its own without any external influence.  And the thermodynamic quantity that we use to describe a spontaneous reaction is free  energy.  But if the change in free energy for that reaction is negative, we know that it's spontaneous.  So we're looking at a spontaneous reaction and this is written in the form of the general  reaction or the overall reaction.  Now in many textbooks they use the term molecular equation, but as all of you can see, we don't  have molecules involved here.  We have ionic compounds, alright?  Therefore I'm just going to call it the general form of the reaction, okay?  Now if we wanted to write the complete ionic equation, then we would write out all the  dissociated species.  So if you start with zinc, that's metallic zinc, so that does not dissociate in water.  But we know copper sulfate is a soluble salt.  How do you know?  Solubility rules, okay?  The solubility rules tells you that copper sulfate is soluble and if it is a soluble  salt, we know that it dissociates completely to give me copper 2 plus aqueous plus SO4  2 minus aqueous.  So in reality this is what the reactants look like in solution.  Now on the product side, now we have zinc sulfate which is a soluble salt.  So the zinc 2 plus and SO4 2 minus would be dissociated and then you end up with copper  metal.  So that would be the complete ionic equation.  And if you look at the complete ionic equation, you can see there's sulfate at the beginning,  there's sulfate at the end.  So in this reaction, the sulfate anions are the spectator ions.  And now if we cancel out the spectator ions on either side, then we end up with the net  ionic equation.  And that would be zinc solid plus copper 2 plus aqueous, giving me zinc 2 plus aqueous  plus copper solid.  So now you can see that this, the net ionic equation actually describes the net event  that takes place.  So this tells you exactly what's going on.  So the net ionic equation kind of summarizes exactly what's going on.  And if you look at this, one way to figure that out is to assign oxidation numbers.  So this is a metal in its elemental state, that would be zero.  This has a plus 2 oxidation state.  This has a plus 2 oxidation state and this is copper metal.  So you can see that if I take zinc, the oxidation number of zinc goes from zero to plus 2.  And we know that when the oxidation number goes up, what is that process called?  Oxidation.  And if you look at the relative change, it changes by 2.  And since oxidation represents loss of electrons, we know that two electrons are lost because  the oxidation number changes by 2.  All right?  Now let's look at copper.  If you look at copper, it goes from plus 2 to zero.  So copper, the oxidation number decreases.  And if it decreases, we know it's a reduction.  And it changes by a magnitude of 2.  And therefore, we know two electrons are gained.  So two electrons are lost and two electrons are gained.  And therefore, you know that two electrons are being transferred in this reaction.  And if you wanted to write this as half reactions, remember we said half reactions don't really  take place.  It's just a convenient way for us to keep track of what's going on.  And a half reaction is purely a mental process for us to kind of figure out what's going  on.  Okay?  So if you want to write half reactions, then the oxidation half reaction would be where  zinc solid goes to zinc 2 plus aqueous plus two electrons.  And this is the oxidation half reaction.  All right?  It's oxidation because the electrons on the product side, electrons are being lost.  Our electrons are being produced here.  So that will be the oxidation.  And then on the other side, we have copper 2 plus aqueous picking up the two electrons  that are lost to give you copper solid.  And this would be the reduction half reaction.  And in this case, the electrons are on the reactant side.  That means they're being picked up in the reaction.  So electrons are being gained in the process.  And therefore, that is the reduction.  Okay?  Now just like in acid-base reaction, remember acid-base reactions involve proton transfer.  Redox reactions involve electron transfer.  Just like a proton is lost and a proton is gained in an acid-base reaction, in a redox  reaction electrons are lost and electrons are gained.  So just like in acid-base reactions, remember we talked about conjugate acid-base pairs  and we always write down conjugate acid-base pairs.  Likewise, we can write redox couples.  So just like conjugate acid-base pairs, I used to describe acid-base reactions.  Here we use the term redox couple.  And in the redox couple, you always write the oxidized form first, then a diagonal line,  and then the reduced form.  So always you write it in that order.  So in this reaction, you can see there are two redox couples.  All right?  One is zinc and the other is copper.  You always write the oxidized form first, so that would be zinc 2 plus, which has oxidation  number of 2, and then zinc solid.  The physical state would be aqueous.  And then on the other side, we have copper 2 plus aqueous and copper solid.  So in this instance, you would always write the redox couple.  And you can see if you have an oxidation reduction reaction, you will always have two pairs of  the redox couples.  All right?  So the redox couples here are always written in the form of the oxidized form, a diagonal  line, and then the reduced form.  Is that clear?  So if any reaction you should be able to identify the redox couples, just like we did for conjugate  acid-base pairs.  Now in the case of conjugate acid-base pairs, it's easy because the acid form has an extra  proton.  The basic form has lost that proton.  So you can just look at that and see the proton lost, and you can identify which is  the acid and which is the conjugate base.  Here it's not directly apparent.  You have to look at oxidation numbers to figure out which is the oxidized form and then which  is the reduced form.  All right?  And write it in that form.  Now the reason we're looking at all of this is ultimately in electrochemistry, what we're  interested in looking at are called electrochemical cells.  So we're going to start looking at electrochemical cells.  And let's define what they are.  So ultimately we're going to look at what we call galvanic cells, or sometimes the term  voltaic cells are used.  And both the galvanic cells or the voltaic cell is called electrochemical cell.  So let me define what an electrochemical cell is.  An electrochemical cell is a device in which an electric current, all right?  Now all of you know electric current is what?  A flow of electrons.  A flow of electrons produced by a spontaneous chemical reaction or is used to bring about  a non-spontaneous reaction.  I know what that buzzing sound is.  I know it's annoying, but it's a whole annoying.  So let's go back.  So we're looking at galvanic cells.  And galvanic cells, voltaic cells are called electrochemical cells.  Overall electrochemical cells are devices that we make up, we build.  And they can be a few types.  One is that you take advantage of a spontaneous reaction to generate a flow of current, all  right?  Or you can take an external source of current.  So you plug it into a plug, for example, that's an external source of current and you use  that to drive a non-spontaneous reaction.  So essentially we're interested in looking at spontaneous reactions and non-spontaneous  reactions, all right?  We're going to start by looking at those that take advantage of a spontaneous reaction.  And one that takes advantage of a spontaneous reaction is called a galvanic cell.  Okay, so a galvanic cell, for voltaic cells, is an electrochemical cell in which a spontaneous  chemical reaction is used to generate an electric current.  All right, so what we're going to do is we're going to start by looking at spontaneous reactions.  And so if you take a galvanic cell or a voltaic cell, this is an electrochemical device.  It's a device that you construct and it takes advantage of a spontaneous chemical reaction  to generate electric current.  So bottom line, what are galvanic cells or voltaic cells?  So these are batteries.  So in electrochemistry we call them galvanic cells or voltaic cells and the simple idea  behind this is that if you want to construct a battery and you know that there are a whole  host of batteries out there, they're lithium batteries, there are different kinds of batteries  that are used for different purposes.  Now all of these are galvanic cells.  The underlying principle is the same and they take advantage of a spontaneous reaction.  Okay?  So let's take an example.  So we just looked at the fact that if you take this reaction where you have zinc solid  and copper 2 plus aqueous, we know that goes to zinc 2 plus aqueous plus copper solid and  this is spontaneous.  All right?  And that means delta G is negative.  So we know that if I take a chunk of zinc metal, all right, and take a solution of copper  sulphate, copper sulphate is a soluble salt.  So you take a spoonful of copper sulphate, put it in water and stir it up and then put  a chunk of zinc metal inside.  What's going to happen is this is a spontaneous reaction.  There's a natural tendency for this reaction to take place and you're going to have the  zinc metal dissolving as it reacts and then at the bottom you're going to see copper metal  depositing at the bottom, all right, as this reaction progresses.  And at the end it's going to establish equilibrium, all right?  Every reaction, remember, ultimately goes to equilibrium.  So there comes a point in which the reaction will start slowing down, not really slowing  down, but the amounts of the reactants and products start decreasing and there comes  a point where the reactants start decreasing, the products are going up and there comes  a point when the amounts of both stop changing and at that point we know the equilibrium  is established, okay?  Now when we analyze this reaction we know two electrons are being transferred, all right?  So that means if I have one atom of zinc and one copper ion, we know two electrons will  be transferred, all right?  So if you just have one atom and one ion of each, you're going to have two electrons transferred.  Now if I have billions upon billions upon billions of zinc atoms and billions upon billions  upon billions of copper atoms, now how many electrons will be transferred?  Twice as much billions, so twice as much billions upon billions of electrons will be  transferred.  Now if I take one mole of zinc and I have one mole of copper, we know two moles of electrons  will be transferred.  You see that?  If you have amounts of both, you're going to have a tremendous amount of electrons being  transferred, okay?  Now the idea behind batteries is quite clever.  The idea is that if you take just the two of them and the intimate contact with each  other, the electrons are going to be transferred and the system will go to equilibrium.  Now what if we could separate the systems so that now we're going to take advantage  of this spontaneous reaction, but if we can physically separate this from this, all right?  And then if we connect them through a wire, can we force these billions upon billions  of electrons, can we force it through the wire?  And in that way, you can now generate a current, all right?  So the idea is instead of having the two reactants in intimate contact with each other, where  the electrons will be transferred from one to the other, rather we want to construct  a device where we're going to take advantage of the fact that there is a spontaneous reaction,  that electrons are going to go from one to the other, but now we're going to force these  electrons through a wire, all right?  Physically separate them and force it through a wire so that now you have electron flow  through the wire.  Got it?  And remember, electron flow is what?  You can get a current, okay?  So that means we have to construct a device where you have all the elements for this process  to take place.  So we have to separate the zinc from the copper.  In order to separate that, what we deal with is where we have the zinc metal, and we're  going to call the zinc metal, we're going to call this the electrode.  And this zinc metal is going to be dipped in a solution of zinc sulfate, okay?  Usually this is about one molar concentration, and so since this is soluble, what you'll end  up with is zinc 2 plus equus and sulfate 2 minus equus, each of which will be one molar.  So on one side, you have the reactant zinc, and so that's one half of that redox reaction,  okay?  On the other side, you're going to have copper, so this will be copper metal, and over here  I'm going to have zinc metal.  This drawing is really bad, but I'll show you a nice picture later, okay?  And this is dipped in copper sulfate.  So this is the copper electrode, and we have copper sulfate equus, AQ, and this will be  one molar typically, and this dissociates to give you copper equus and sulfate because  it's a soluble salt, okay?  Now I'm going to link these through a wire, and if you wanted to just to make sure that  there's a current flowing through, we can link it through a voltmeter.  So this is a voltmeter, alright?  And so now I've physically separated the oxidation half reaction from the reduction half reaction  so that we have the two reactants separated, and we're going to force the electrons to  that wire, but if I'm going to force the electrons to flow to that wire, all of you  recognize that I have to complete the circuit, okay?  The circuits are completed.  In order to complete the circuit, what we have is what we call a salt bridge.  So this is a salt bridge, and a salt bridge essentially is a gel that's in that U-shaped  glass tube, and over here I'm going to put a cotton plug to prevent the gel from seeping  out, alright?  And in here, this gel has KCl equus.  In other words, it has potassium ions equus plus Cl minus equus, alright?  And I'm now going to switch and show you a pretty picture because it's the same thing  that I drew, alright?  So now I'm going to switch to a picture that kind of shows you.  So let's start with this picture in the middle.  So we have the two electrodes.  So we have the zinc electrode and zinc metal dipped in zinc sulfate, and that dissociates  to give you zinc ions and sulfate ions.  On the other side, we have separated the other reactant, which is copper, and that is copper  metal dipped in copper sulfate, and that has copper ions and sulfate ions, and the two  are connected through a wire, and we're going to attach it to a voltmeter so that we can  measure to see where there's a flow of current or not, okay?  Now in order to complete the circuit, what we have is a salt bridge.  So here is the salt bridge, and the salt bridge is essentially a gel that's placed in there  that has ions in it, dissociated ions, and so you have potassium ions and chloride ions  dissociated in that, and then at the bottom of the glass YouTube that has the gel, you  have cotton plugs on either side just so that your gel doesn't come flopping out, all right?  Now when we complete the circuit, we know that when you have flow of electrons, you  have a current flowing.  Another, a different way in which you can have current flowing is also to have motion  of ions, all right?  If you have movement of ions too, you can have a flow of current, all right?  And so what's happening here is that you have a flow of electrons that go from this to this,  but now you complete the circuit because now you can have ions flowing from this to this,  all right?  And so overall what we've done is we've completed the circuit.  So there's a complete circuit so that now you can have electrons flowing from one side  to the other, got it?  So this is simply put, this is a very rudimentary battery.  This is the idea behind batteries, all right?  Now if you look at this, at the left hand side is where the oxidation takes place.  So in this, on the left hand side, this is where, we call this the anode, and the anode  is where the oxidation occurs, okay?  And the sign for the anode, the anode has a negative sign, all right?  And what happens at the anode is electrons are pushed into the wire, all right?  So if you look at the left hand side, usually it's convenient to draw these out so that  you put it in alphabetical order.  One side is anode, the other side is called a cathode, and therefore it always comes first,  all right?  So at the anode, you always have oxidation taking place, and when oxidation occurs, what  happens?  You have electrons being produced, it's electron loss.  So when the oxidation takes place, you have lots of electrons being produced, or lots  of electrons being lost.  Remember those electrons are going to be channeled through the wire.  So at the anode, electrons are being pushed into the wire so that you have a flow of electrons,  okay?  Now on the other side, you have the cathode.  So and this is where reduction occurs, and this has, is symbolized by a positive sign,  and we say electrons are pulled in from the wire, all right?  So the wire is connected to the copper electrode, and so the cathode is where electrons are  being pulled in.  So any electrochemical or galvanic cell has two components.  If you take a galvanic cell, it's got an anode, because we've separated the two reactions.  If they were in contact with each other, now the electrons will be transferred one to the  other, if they're in contact with each other, then that's not useful to us.  So we're taking advantage of a spontaneous reaction that we know has a natural tendency  to occur without any external influence.  We're going to physically separate the oxidation half reaction from the reduction half reaction,  and we're going to connect the two through a wire so that the spontaneous reaction still  occurs.  So now you're forcing those electrons, those billions upon billions upon billions of electrons,  now are being channeled through that wire from one to the other, all right?  So wherever you have the oxidation, we call that the anode.  The anode always has a negative sign, and then you know that at the anode, electrons  are being produced, and those electrons are being forced through the wire.  They're going to push through the wire.  And at the other end, the flow of electrons comes into the cathode.  So the cathode is where the reduction takes place, where electrons are being gained, and  so the wire, the electrons are taken up, and that's where the reduction takes place, and  so you have a pull of electrons.  So simple put, what you have here is a push and pull of electrons.  You have one terminal, you're having electrons being pushed to the wire, and at the other  terminal, you're having electrons being pulled in, okay?  And I kind of like this picture because it summarizes everything that's going on.  So if you look at this, and if you look at the bottom here, you can see that this is  where the oxidation takes place, and zinc is oxidized to zinc 2 plus.  So the net reaction is zinc is going to zinc 2 plus plus 2 electrons.  At the oxidation point, two electrons are being produced.  If you have the ability to magnify what's going on, this is the zinc electrode, which  is a metallic solid, all right?  So the zinc electrode will give off two electrons, and it becomes zinc 2 ions, and when it becomes  an ion, it dissolves in solution.  You guys see that?  So what's happening at the anode is that the metal is actually dissolving because the metal  is giving up, the metallic solid is giving up two electrons and pushing it up through  the wire and dissolving as zinc 2 plus.  So at the anode, what's going to happen is if you start with a solid metal strip of zinc  and you weigh it before you start, you connect everything up and complete the circuit, if  you weighed before and then you weighed afterwards, what's going to happen to the mass of that  strip, metal strip?  It's going to go down.  Can everybody see that?  Because at the anode, oxidation takes place, and that means the metallic solid is now  converting to zinc 2 plus and the zinc 2 plus dissolves in solution.  All right?  Now if you go to the other side where the reduction takes place, now here's the reduction half  reaction.  So the electrons that come into the copper electrode are being pulled in, and the two  electrons react with copper 2 plus, the copper 2 plus ions in solution, and you form a copper  solid.  So what happens at the cathode is you're having more and more copper depositing on that electrode.  So if you weigh that electrode before anything happens, and if you weigh it afterwards, you'll  see that its mass goes up.  All right?  And so what's going to happen is at the anode, your electrode will start dissolving and getting  smaller and smaller and smaller, whereas here where you have the dissolved copper ions picking  up the two electrons.  So they're going to pick up the two electrons to give you copper solid, and these will deposit  on the electrode, and therefore the mass of the electrode increases.  All right?  And this will continue until the reaction reaches equilibrium.  And when the reaction reaches equilibrium, now you don't have any net flow of electrons.  All right?  And that's what we mean when we, in real life, we know that all batteries go dead after a  while.  And when a battery goes dead, what does that mean?  That equilibrium has been established.  All right?  Once equilibrium is established, now delta G equals zero.  All right?  And therefore there's no flow of electrons, and so the battery goes dead.  How do we know that there is an electron flow?  Use a voltmeter to measure it.  Now, have you guys done a rudimentary electrochemical cell in the lab yet?  This week.  Okay.  So now, one of the things in the lab is if you attach, you set up your electrochemical  cell, you know what your anode is where the oxidation is.  You set up your cathode, you know where the reduction takes place, and you hook it up  to your voltmeter, and your voltmeter reads a negative sign.  That means that you reverse the terminal, because when you measure the voltage, the voltage  should always be positive.  All right?  And so when you hook it up to the voltmeter and complete the circuit, and if you see a  negative sign on your voltmeter, that means you reverse the terminals.  That means you, you know, remember the anode is the negative terminal, the cathode is  the positive terminal, and therefore you have to hook it up appropriately.  Got it?  Now we said lastly that you have to complete the circuit, and we're going to use the movement  of cations and anions to complete the circuit.  All right?  And if you look at this, what happens is that not only are you completing the circuit and  this movement of cations and anions help the flow of current through the system, but also  they also help maintain charge balance.  So now if you look at the zinc side, remember zinc solid is going to zinc 2 plus.  So you're starting with something that's electrically neutral, and in that solution you're having  more and more cations.  Remember everything in nature has to be electrically balanced.  So if you're forming more and more cations, you need anions to counterbalance that charge.  And so what happens is in this gel, to counterbalance the extra zinc ions that are coming in, the  anions start flowing into that solution to balance the charge.  So you have chloride ions moving in, migrating through that gel, they'll migrate and balance  the charge as well.  So not only do the flow of the ions generate a current, and you can complete that circuit,  but also they help maintain charge balance.  Now on the other side, can you see that you have copper ions depositing?  So you have copper 2 plus cations going to something that's electrically neutral.  So in this solution, if all of your cations become neutral, you can see you have lots  of anions in there that don't have a counter charge to balance that charge.  So to balance, because in that solution you have lots of SO4 2 minus, the sulfate anions  are there, only the cation becomes electrically neutral, and therefore you need cations to  counterbalance those sulfate ions.  And so to balance the charge, you have migration of potassium ions that migrate through the  gel into that.  Got it?  And so that describes a galvanic cell.  Yes, go ahead.  The sulfate is a specter ion, okay?  But you need sulfates because everything in nature has to be electrically neutral.  So if you have a cation, a cation always has to have an anion that balances the charge.  Do you see that?  To begin with, before any reaction takes place, you have the same number of zinc ions and the  same number of sulfate ions in here, and you have the same number in this solution, the  initial amount of copper ions will equal the sulfate ions.  Got it?  But once the reaction starts proceeding, you know that your zinc is, your zinc metal  becomes zinc 2 plus, so in this solution, you're going to have more and more cations  building up.  So to balance the charge of the increasing number of cations, you need anions to migrate  in there to balance the charge.  Is that again?  You have to have the sulfate.  In nature, you can't have cations without anions.  Do you understand that?  So take sodium chloride.  You can't have, if you take sodium chloride, which is an ionic salt, you cannot have sodium  cations alone.  Did you understand that?  So you need chlorides to balance that.  Yes?  Is that again?  You can't have a salt bridge.  Salt bridge.  Salt bridge.  Without the salt bridge.  Is that what your question was?  Okay.  Without the salt bridge, the circuit won't be complete.  Do you understand that?  You need the salt bridge because for electrical flow, your electrical circuit has to be completed.  All right?  And so if you didn't have the salt bridge, nothing would happen because you have to have  the circuit completed.  Okay?  Sorry.  I miss, I heard the sulfate, not salt bridge.  Okay.  So we've looked at an example and this particular cell is called the galvanic cell.  The galvanic cell has a particular name and this is called the danial cell.  So if, so we call this the danial cell, so let me just, so if your reactants are zinc,  metal and you have zinc, zinc sulfate and copper, copper sulfate, this combination of  a galvanic cell has a specific name and this is called the danial cell.  All right?  Now you can have different types of electrodes.  So let's take another example.  So if you want to describe another example of a cell, let's take another electrode.  So we looked at electrodes and we've looked at zinc electrode.  We've looked at copper electrodes.  All right?  So in the danial cell, the zinc electrode forms the anode, the copper electrode forms  the cathode and you have a flow of electrons.  Okay?  Now you can also have another electrode called the hydrogen electrode.  So let's take a look at what it means when we say we have a hydrogen electrode.  And let's say your hydrogen electrode is the cathode.  So if your hydrogen electrode is the cathode, remember that the cathode, does oxidation  take place or reduction take place?  Reduction.  All right?  So the reduction reaction that you see at the cathode will be 2H plus equus plus 2 electrons  giving you hydrogen gas.  All right?  And so if you wanted to set this up, your H plus comes in the solution in the form of  let's say HCl equus and usually the concentration is one molar.  So you can see in this, it's a strong acid, it dissociates completely.  So in that solution you have H plus ions.  All right?  Now the product is hydrogen.  So we need something.  Remember we need a wire to bring in the electrons into, remember where the reduction takes place,  electrons are being pulled in.  So you need a flow of electrons coming into that solution.  And so what we do is we have an inert metal like that and this whole, the wire as well  as this is platinum and this is an inert metal.  So the reason it's inert is it doesn't react.  All it does is just serve to deliver the electrons into the solution and you have a test tube  that looks something like this.  So you have a test tube which is open at one end like that and you have hydrogen gas as  well.  And so essentially what happens is the hydrogen gas comes through here and then it'll bubble  out on these sides.  You'll see bubbles coming out and this would be the cathode.  All right?  So if you wanted to look at a setup where your cathode was now a hydrogen electrode, this  is what it would look like.  So here on this side we have the anode and the anode is the zinc electrode.  The cathode is a hydrogen electrode.  So essentially what you do is you have hydrogen gas at one atmosphere, you have one molar  HCl, your platinum, the platinum metal is just used because it's an inert metal.  You need something to deliver the flow of electrons into that solution so that this reduction  occurs and so an inert material like platinum is used.  All right?  And you can see the bubbles of hydrogen gas.  So if you wanted to describe what's going on in this cell then you would say that at  the cathode, so at the anode you would have the oxidation taking place which is the zinc  solid going to zinc 2 plus equus plus 2 electrons.  At the cathode you would have the 2H plus ions picking up 2 electrons to give you H2  gas.  And so your net ionic equation would be where you have zinc solid plus 2H plus equus giving  you zinc 2 plus equus plus H2 gas.  Did you guys get that?  So overall the net ionic equation that takes place in this galvanic cell is shown here and  you've identified what the anode is, you've identified what the cathode is and then how  many electrons are being transferred?  There are 2 electrons being transferred in that setup.  Now you could also have the other way around.  So now let's say let's take a different cell and in the other cell let's say the oxidation.  So what we have is the anode is where you have the oxidation.  So now let's say we have the hydrogen electrode as the anode.  So the hydrogen electrode is the anode, you're going to have hydrogen giving you 2H plus  equus plus 2 electrons.  This would be the oxidation.  At the cathode you would have copper 2 plus equus pick up the 2 electrons to give you  copper solid.  So this is the reduction.  And now we've looked at an example where the hydrogen electrode is the anode and now the  copper electrode is the cathode.  So we took the Daniel cell and we took one half of it and replaced one half with the  hydrogen electrode on either side.  And you can see that the net reaction would be, so the net ionic equation would be where  you have hydrogen gas plus copper 2 plus equus giving you 2H plus equus plus copper solid.  So the overall reaction that takes place in the second example where now your hydrogen  electrode is the anode where the oxidation takes place and now copper is the cathode  then the overall reaction would look like that.  So that's what this picture shows you.  So the two examples that we looked at here is remember usually when you draw these out  you're customary to always put the anode first and then the cathode second.  Here you have the zinc electrode here and over here you have the hydrogen electrode.  So this is the cathode and you have a flow of electrons and the circuit is completed with  the salt bridge.  Now over here this is the other.  Now you have a copper electrode which is the cathode.  Your hydrogen electrode is now the anode so now this is where the oxidation occurs.  This is where electrons are being pushed out and it goes around and the electrons on this  side are being pulled in.  You see that?  These are all examples of galvanic cells.  A galvanic cell is an electrochemical device that takes advantage of a spontaneous reaction.  It has to be a spontaneous reaction if you're going to have a flow of electrons to the wire.  And so galvanic cell is another name for batteries.  This is the underlying principle for all batteries.  So now we're going to look at the last few minutes.  We're going to look at what we call cell notations.  So this is chemist love shorthand notations.  And so we have developed this system of notations that we use to describe cells, particularly  galvanic cells.  So we're going to look at cell notations.  And let's start with the Daniel cell.  We know in the Daniel cell, the reaction that we're looking at, the net reaction, is that  you have zinc solid plus copper 2 plus ions giving you zinc 2 plus equus plus copper solid.  So that isn't it.  And so in this reaction, we looked at the fact that the redox couples are, we have the  oxidized form first, so the zinc 2 plus equus, we use a diagonal line and then zinc solid.  And we have copper 2 plus equus, copper solid.  That's the redox couple.  All right?  Now if we want to describe the cathode for this reaction, let's start with the anode because  always in alphabetical order.  So let's start with anode.  Now when you write the anode, and this is called the notation for the electrode.  So this is, we're looking at the electrode.  You always write the reactant first and then the product.  All right?  And if there's a vertical line like that, now it's different from a diagonal line.  It's very, very important that you draw vertical lines and diagonal lines and don't mix up  the two.  All right?  If you draw a vertical line like that, it means that there, the reactants and products  are in two different phases.  So if you have something that's dissolved in aqueous solution and a solid, or if you  have aqueous and a gas, or you have a gas and a liquid, all right?  These are two different phases and what we say is that there's an interface between them.  All right?  So if you have a vertical line, there are two different phases.  Or if you always write the reactant first and then you put a comma and then product,  this means same phase.  All right?  So if I wanted to write the Daniel cell, my anode would be, my reactant for that is zinc.  I'm going to put a vertical line because we're going from zinc to zinc 2 plus aqueous.  All right?  So that describes the anode.  Now if I want to describe the cathode, I would write the reactant.  Vertical line because again, you're going from an aqueous solution to a solid so there's  a change in phase would be that.  All right?  Now we describe the cell, so the cell notation is always anode first, cathode last.  All right?  So just remember it as AC.  All right?  It's an alphabetical order.  Just remember that.  Anode always comes first.  All right?  And therefore, I would start with the anode.  Remember with the redox couple, it's always the oxidized form, diagonal sign, reduced form.  When you draw the symbols for the electrodes, the notation, it's always reactant first,  product.  And if there's a phase change, you put a vertical line between them.  Okay?  So now if I combine them to give the cell notation, what you do is you start with anode  and you know the anode is that, so vertical line, zinc 2 plus aqueous.  Now I put a double line and this double line refers to the salt bridge which completes  the circuit.  And then on the other side, we have the cathode reactant first vertical line product.  So what does this symbol convey?  To chemists, this symbol represents this picture.  So every time you see this, that's what this picture represents.  So the shorthand notation is a very convenient way to convey to everybody.  That symbol, whenever you see those sets of symbol, you have to have a mental picture  of what your galvanic cell looks like.  And so that translates to this.  So you know that the anode has a zinc electrode dipped in zinc 2 plus, alright?  And that it has a salt bridge and on the other side, you have a copper dipped in copper  2 plus and it signifies this picture.  Got it?  Now at the beginning of next class, I will take a few more examples of writing the cell  notation and show you how to write cell notations to describe what the galvanic cell looks like.  Okay?  So I will stop there for today.
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Okay, so were there any questions about the material from last week? Yes, go ahead. The lecture from Friday wasn't posted yet. Okay, it should be up, so let me check and let you guys know. So the lecture from Friday, I think, is up on the website, isn't it, Sean? Okay, it is up. Can you check and let me know? Send me an email and I'll forward it to him. Okay? Anything else? Yes? If you're an acid, yes. If you're balancing redox reactions and acid acidic solution, it's a lot easier to put H plus rather than H3O plus. Alright? But technically speaking, it's more accurate to put H3O plus, okay? But it's a lot easier to balance it if you just have H plus. Got it? Okay, so today we're going to move on to the actual topic which deals with electrochemistry. And we said electrochemistry deals with redox reactions. And so let's take an example of a redox reaction, the one written here where you have zinc metal reacting with copper sulfate to give you zinc sulfate and copper. Now this is a spontaneous reaction. Remember last quarter when we looked at thermodynamics, when a reaction is spontaneous, it means that it has a natural tendency to occur on its own without any external influence. And the thermodynamic quantity that we use to describe a spontaneous reaction is free energy. But if the change in free energy for that reaction is negative, we know that it's spontaneous. So we're looking at a spontaneous reaction and this is written in the form of the general reaction or the overall reaction. Now in many textbooks they use the term molecular equation, but as all of you can see, we don't have molecules involved here. We have ionic compounds, alright? Therefore I'm just going to call it the general form of the reaction, okay? Now if we wanted to write the complete ionic equation, then we would write out all the dissociated species. So if you start with zinc, that's metallic zinc, so that does not dissociate in water. But we know copper sulfate is a soluble salt. How do you know? Solubility rules, okay? The solubility rules tells you that copper sulfate is soluble and if it is a soluble salt, we know that it dissociates completely to give me copper 2 plus aqueous plus SO4 2 minus aqueous. So in reality this is what the reactants look like in solution. Now on the product side, now we have zinc sulfate which is a soluble salt. So the zinc 2 plus and SO4 2 minus would be dissociated and then you end up with copper metal. So that would be the complete ionic equation. And if you look at the complete ionic equation, you can see there's sulfate at the beginning, there's sulfate at the end. So in this reaction, the sulfate anions are the spectator ions. And now if we cancel out the spectator ions on either side, then we end up with the net ionic equation. And that would be zinc solid plus copper 2 plus aqueous, giving me zinc 2 plus aqueous plus copper solid. So now you can see that this, the net ionic equation actually describes the net event that takes place. So this tells you exactly what's going on. So the net ionic equation kind of summarizes exactly what's going on. And if you look at this, one way to figure that out is to assign oxidation numbers. So this is a metal in its elemental state, that would be zero. This has a plus 2 oxidation state. This has a plus 2 oxidation state and this is copper metal. So you can see that if I take zinc, the oxidation number of zinc goes from zero to plus 2. And we know that when the oxidation number goes up, what is that process called? Oxidation. And if you look at the relative change, it changes by 2. And since oxidation represents loss of electrons, we know that two electrons are lost because the oxidation number changes by 2. All right? Now let's look at copper. If you look at copper, it goes from plus 2 to zero. So copper, the oxidation number decreases. And if it decreases, we know it's a reduction. And it changes by a magnitude of 2. And therefore, we know two electrons are gained. So two electrons are lost and two electrons are gained. And therefore, you know that two electrons are being transferred in this reaction. And if you wanted to write this as half reactions, remember we said half reactions don't really take place. It's just a convenient way for us to keep track of what's going on. And a half reaction is purely a mental process for us to kind of figure out what's going on. Okay? So if you want to write half reactions, then the oxidation half reaction would be where zinc solid goes to zinc 2 plus aqueous plus two electrons. And this is the oxidation half reaction. All right? It's oxidation because the electrons on the product side, electrons are being lost. Our electrons are being produced here. So that will be the oxidation. And then on the other side, we have copper 2 plus aqueous picking up the two electrons that are lost to give you copper solid. And this would be the reduction half reaction. And in this case, the electrons are on the reactant side. That means they're being picked up in the reaction. So electrons are being gained in the process. And therefore, that is the reduction. Okay? Now just like in acid-base reaction, remember acid-base reactions involve proton transfer. Redox reactions involve electron transfer. Just like a proton is lost and a proton is gained in an acid-base reaction, in a redox reaction electrons are lost and electrons are gained. So just like in acid-base reactions, remember we talked about conjugate acid-base pairs and we always write down conjugate acid-base pairs. Likewise, we can write redox couples. So just like conjugate acid-base pairs, I used to describe acid-base reactions. Here we use the term redox couple. And in the redox couple, you always write the oxidized form first, then a diagonal line, and then the reduced form. So always you write it in that order. So in this reaction, you can see there are two redox couples. All right? One is zinc and the other is copper. You always write the oxidized form first, so that would be zinc 2 plus, which has oxidation number of 2, and then zinc solid. The physical state would be aqueous. And then on the other side, we have copper 2 plus aqueous and copper solid. So in this instance, you would always write the redox couple. And you can see if you have an oxidation reduction reaction, you will always have two pairs of the redox couples. All right? So the redox couples here are always written in the form of the oxidized form, a diagonal line, and then the reduced form. Is that clear? So if any reaction you should be able to identify the redox couples, just like we did for conjugate acid-base pairs. Now in the case of conjugate acid-base pairs, it's easy because the acid form has an extra proton. The basic form has lost that proton. So you can just look at that and see the proton lost, and you can identify which is the acid and which is the conjugate base. Here it's not directly apparent. You have to look at oxidation numbers to figure out which is the oxidized form and then which is the reduced form. All right? And write it in that form. Now the reason we're looking at all of this is ultimately in electrochemistry, what we're interested in looking at are called electrochemical cells. So we're going to start looking at electrochemical cells. And let's define what they are. So ultimately we're going to look at what we call galvanic cells, or sometimes the term voltaic cells are used. And both the galvanic cells or the voltaic cell is called electrochemical cell. So let me define what an electrochemical cell is. An electrochemical cell is a device in which an electric current, all right? Now all of you know electric current is what? A flow of electrons. A flow of electrons produced by a spontaneous chemical reaction or is used to bring about a non-spontaneous reaction. I know what that buzzing sound is. I know it's annoying, but it's a whole annoying. So let's go back. So we're looking at galvanic cells. And galvanic cells, voltaic cells are called electrochemical cells. Overall electrochemical cells are devices that we make up, we build. And they can be a few types. One is that you take advantage of a spontaneous reaction to generate a flow of current, all right? Or you can take an external source of current. So you plug it into a plug, for example, that's an external source of current and you use that to drive a non-spontaneous reaction. So essentially we're interested in looking at spontaneous reactions and non-spontaneous reactions, all right? We're going to start by looking at those that take advantage of a spontaneous reaction. And one that takes advantage of a spontaneous reaction is called a galvanic cell. Okay, so a galvanic cell, for voltaic cells, is an electrochemical cell in which a spontaneous chemical reaction is used to generate an electric current. All right, so what we're going to do is we're going to start by looking at spontaneous reactions. And so if you take a galvanic cell or a voltaic cell, this is an electrochemical device. It's a device that you construct and it takes advantage of a spontaneous chemical reaction to generate electric current. So bottom line, what are galvanic cells or voltaic cells? So these are batteries. So in electrochemistry we call them galvanic cells or voltaic cells and the simple idea behind this is that if you want to construct a battery and you know that there are a whole host of batteries out there, they're lithium batteries, there are different kinds of batteries that are used for different purposes. Now all of these are galvanic cells. The underlying principle is the same and they take advantage of a spontaneous reaction. Okay? So let's take an example. So we just looked at the fact that if you take this reaction where you have zinc solid and copper 2 plus aqueous, we know that goes to zinc 2 plus aqueous plus copper solid and this is spontaneous. All right? And that means delta G is negative. So we know that if I take a chunk of zinc metal, all right, and take a solution of copper sulphate, copper sulphate is a soluble salt. So you take a spoonful of copper sulphate, put it in water and stir it up and then put a chunk of zinc metal inside. What's going to happen is this is a spontaneous reaction. There's a natural tendency for this reaction to take place and you're going to have the zinc metal dissolving as it reacts and then at the bottom you're going to see copper metal depositing at the bottom, all right, as this reaction progresses. And at the end it's going to establish equilibrium, all right? Every reaction, remember, ultimately goes to equilibrium. So there comes a point in which the reaction will start slowing down, not really slowing down, but the amounts of the reactants and products start decreasing and there comes a point where the reactants start decreasing, the products are going up and there comes a point when the amounts of both stop changing and at that point we know the equilibrium is established, okay? Now when we analyze this reaction we know two electrons are being transferred, all right? So that means if I have one atom of zinc and one copper ion, we know two electrons will be transferred, all right? So if you just have one atom and one ion of each, you're going to have two electrons transferred. Now if I have billions upon billions upon billions of zinc atoms and billions upon billions upon billions of copper atoms, now how many electrons will be transferred? Twice as much billions, so twice as much billions upon billions of electrons will be transferred. Now if I take one mole of zinc and I have one mole of copper, we know two moles of electrons will be transferred. You see that? If you have amounts of both, you're going to have a tremendous amount of electrons being transferred, okay? Now the idea behind batteries is quite clever. The idea is that if you take just the two of them and the intimate contact with each other, the electrons are going to be transferred and the system will go to equilibrium. Now what if we could separate the systems so that now we're going to take advantage of this spontaneous reaction, but if we can physically separate this from this, all right? And then if we connect them through a wire, can we force these billions upon billions of electrons, can we force it through the wire? And in that way, you can now generate a current, all right? So the idea is instead of having the two reactants in intimate contact with each other, where the electrons will be transferred from one to the other, rather we want to construct a device where we're going to take advantage of the fact that there is a spontaneous reaction, that electrons are going to go from one to the other, but now we're going to force these electrons through a wire, all right? Physically separate them and force it through a wire so that now you have electron flow through the wire. Got it? And remember, electron flow is what? You can get a current, okay? So that means we have to construct a device where you have all the elements for this process to take place. So we have to separate the zinc from the copper. In order to separate that, what we deal with is where we have the zinc metal, and we're going to call the zinc metal, we're going to call this the electrode. And this zinc metal is going to be dipped in a solution of zinc sulfate, okay? Usually this is about one molar concentration, and so since this is soluble, what you'll end up with is zinc 2 plus equus and sulfate 2 minus equus, each of which will be one molar. So on one side, you have the reactant zinc, and so that's one half of that redox reaction, okay? On the other side, you're going to have copper, so this will be copper metal, and over here I'm going to have zinc metal. This drawing is really bad, but I'll show you a nice picture later, okay? And this is dipped in copper sulfate. So this is the copper electrode, and we have copper sulfate equus, AQ, and this will be one molar typically, and this dissociates to give you copper equus and sulfate because it's a soluble salt, okay? Now I'm going to link these through a wire, and if you wanted to just to make sure that there's a current flowing through, we can link it through a voltmeter. So this is a voltmeter, alright? And so now I've physically separated the oxidation half reaction from the reduction half reaction so that we have the two reactants separated, and we're going to force the electrons to that wire, but if I'm going to force the electrons to flow to that wire, all of you recognize that I have to complete the circuit, okay? The circuits are completed. In order to complete the circuit, what we have is what we call a salt bridge. So this is a salt bridge, and a salt bridge essentially is a gel that's in that U-shaped glass tube, and over here I'm going to put a cotton plug to prevent the gel from seeping out, alright? And in here, this gel has KCl equus. In other words, it has potassium ions equus plus Cl minus equus, alright? And I'm now going to switch and show you a pretty picture because it's the same thing that I drew, alright? So now I'm going to switch to a picture that kind of shows you. So let's start with this picture in the middle. So we have the two electrodes. So we have the zinc electrode and zinc metal dipped in zinc sulfate, and that dissociates to give you zinc ions and sulfate ions. On the other side, we have separated the other reactant, which is copper, and that is copper metal dipped in copper sulfate, and that has copper ions and sulfate ions, and the two are connected through a wire, and we're going to attach it to a voltmeter so that we can measure to see where there's a flow of current or not, okay? Now in order to complete the circuit, what we have is a salt bridge. So here is the salt bridge, and the salt bridge is essentially a gel that's placed in there that has ions in it, dissociated ions, and so you have potassium ions and chloride ions dissociated in that, and then at the bottom of the glass YouTube that has the gel, you have cotton plugs on either side just so that your gel doesn't come flopping out, all right? Now when we complete the circuit, we know that when you have flow of electrons, you have a current flowing. Another, a different way in which you can have current flowing is also to have motion of ions, all right? If you have movement of ions too, you can have a flow of current, all right? And so what's happening here is that you have a flow of electrons that go from this to this, but now you complete the circuit because now you can have ions flowing from this to this, all right? And so overall what we've done is we've completed the circuit. So there's a complete circuit so that now you can have electrons flowing from one side to the other, got it? So this is simply put, this is a very rudimentary battery. This is the idea behind batteries, all right? Now if you look at this, at the left hand side is where the oxidation takes place. So in this, on the left hand side, this is where, we call this the anode, and the anode is where the oxidation occurs, okay? And the sign for the anode, the anode has a negative sign, all right? And what happens at the anode is electrons are pushed into the wire, all right? So if you look at the left hand side, usually it's convenient to draw these out so that you put it in alphabetical order. One side is anode, the other side is called a cathode, and therefore it always comes first, all right? So at the anode, you always have oxidation taking place, and when oxidation occurs, what happens? You have electrons being produced, it's electron loss. So when the oxidation takes place, you have lots of electrons being produced, or lots of electrons being lost. Remember those electrons are going to be channeled through the wire. So at the anode, electrons are being pushed into the wire so that you have a flow of electrons, okay? Now on the other side, you have the cathode. So and this is where reduction occurs, and this has, is symbolized by a positive sign, and we say electrons are pulled in from the wire, all right? So the wire is connected to the copper electrode, and so the cathode is where electrons are being pulled in. So any electrochemical or galvanic cell has two components. If you take a galvanic cell, it's got an anode, because we've separated the two reactions. If they were in contact with each other, now the electrons will be transferred one to the other, if they're in contact with each other, then that's not useful to us. So we're taking advantage of a spontaneous reaction that we know has a natural tendency to occur without any external influence. We're going to physically separate the oxidation half reaction from the reduction half reaction, and we're going to connect the two through a wire so that the spontaneous reaction still occurs. So now you're forcing those electrons, those billions upon billions upon billions of electrons, now are being channeled through that wire from one to the other, all right? So wherever you have the oxidation, we call that the anode. The anode always has a negative sign, and then you know that at the anode, electrons are being produced, and those electrons are being forced through the wire. They're going to push through the wire. And at the other end, the flow of electrons comes into the cathode. So the cathode is where the reduction takes place, where electrons are being gained, and so the wire, the electrons are taken up, and that's where the reduction takes place, and so you have a pull of electrons. So simple put, what you have here is a push and pull of electrons. You have one terminal, you're having electrons being pushed to the wire, and at the other terminal, you're having electrons being pulled in, okay? And I kind of like this picture because it summarizes everything that's going on. So if you look at this, and if you look at the bottom here, you can see that this is where the oxidation takes place, and zinc is oxidized to zinc 2 plus. So the net reaction is zinc is going to zinc 2 plus plus 2 electrons. At the oxidation point, two electrons are being produced. If you have the ability to magnify what's going on, this is the zinc electrode, which is a metallic solid, all right? So the zinc electrode will give off two electrons, and it becomes zinc 2 ions, and when it becomes an ion, it dissolves in solution. You guys see that? So what's happening at the anode is that the metal is actually dissolving because the metal is giving up, the metallic solid is giving up two electrons and pushing it up through the wire and dissolving as zinc 2 plus. So at the anode, what's going to happen is if you start with a solid metal strip of zinc and you weigh it before you start, you connect everything up and complete the circuit, if you weighed before and then you weighed afterwards, what's going to happen to the mass of that strip, metal strip? It's going to go down. Can everybody see that? Because at the anode, oxidation takes place, and that means the metallic solid is now converting to zinc 2 plus and the zinc 2 plus dissolves in solution. All right? Now if you go to the other side where the reduction takes place, now here's the reduction half reaction. So the electrons that come into the copper electrode are being pulled in, and the two electrons react with copper 2 plus, the copper 2 plus ions in solution, and you form a copper solid. So what happens at the cathode is you're having more and more copper depositing on that electrode. So if you weigh that electrode before anything happens, and if you weigh it afterwards, you'll see that its mass goes up. All right? And so what's going to happen is at the anode, your electrode will start dissolving and getting smaller and smaller and smaller, whereas here where you have the dissolved copper ions picking up the two electrons. So they're going to pick up the two electrons to give you copper solid, and these will deposit on the electrode, and therefore the mass of the electrode increases. All right? And this will continue until the reaction reaches equilibrium. And when the reaction reaches equilibrium, now you don't have any net flow of electrons. All right? And that's what we mean when we, in real life, we know that all batteries go dead after a while. And when a battery goes dead, what does that mean? That equilibrium has been established. All right? Once equilibrium is established, now delta G equals zero. All right? And therefore there's no flow of electrons, and so the battery goes dead. How do we know that there is an electron flow? Use a voltmeter to measure it. Now, have you guys done a rudimentary electrochemical cell in the lab yet? This week. Okay. So now, one of the things in the lab is if you attach, you set up your electrochemical cell, you know what your anode is where the oxidation is. You set up your cathode, you know where the reduction takes place, and you hook it up to your voltmeter, and your voltmeter reads a negative sign. That means that you reverse the terminal, because when you measure the voltage, the voltage should always be positive. All right? And so when you hook it up to the voltmeter and complete the circuit, and if you see a negative sign on your voltmeter, that means you reverse the terminals. That means you, you know, remember the anode is the negative terminal, the cathode is the positive terminal, and therefore you have to hook it up appropriately. Got it? Now we said lastly that you have to complete the circuit, and we're going to use the movement of cations and anions to complete the circuit. All right? And if you look at this, what happens is that not only are you completing the circuit and this movement of cations and anions help the flow of current through the system, but also they also help maintain charge balance. So now if you look at the zinc side, remember zinc solid is going to zinc 2 plus. So you're starting with something that's electrically neutral, and in that solution you're having more and more cations. Remember everything in nature has to be electrically balanced. So if you're forming more and more cations, you need anions to counterbalance that charge. And so what happens is in this gel, to counterbalance the extra zinc ions that are coming in, the anions start flowing into that solution to balance the charge. So you have chloride ions moving in, migrating through that gel, they'll migrate and balance the charge as well. So not only do the flow of the ions generate a current, and you can complete that circuit, but also they help maintain charge balance. Now on the other side, can you see that you have copper ions depositing? So you have copper 2 plus cations going to something that's electrically neutral. So in this solution, if all of your cations become neutral, you can see you have lots of anions in there that don't have a counter charge to balance that charge. So to balance, because in that solution you have lots of SO4 2 minus, the sulfate anions are there, only the cation becomes electrically neutral, and therefore you need cations to counterbalance those sulfate ions. And so to balance the charge, you have migration of potassium ions that migrate through the gel into that. Got it? And so that describes a galvanic cell. Yes, go ahead. The sulfate is a specter ion, okay? But you need sulfates because everything in nature has to be electrically neutral. So if you have a cation, a cation always has to have an anion that balances the charge. Do you see that? To begin with, before any reaction takes place, you have the same number of zinc ions and the same number of sulfate ions in here, and you have the same number in this solution, the initial amount of copper ions will equal the sulfate ions. Got it? But once the reaction starts proceeding, you know that your zinc is, your zinc metal becomes zinc 2 plus, so in this solution, you're going to have more and more cations building up. So to balance the charge of the increasing number of cations, you need anions to migrate in there to balance the charge. Is that again? You have to have the sulfate. In nature, you can't have cations without anions. Do you understand that? So take sodium chloride. You can't have, if you take sodium chloride, which is an ionic salt, you cannot have sodium cations alone. Did you understand that? So you need chlorides to balance that. Yes? Is that again? You can't have a salt bridge. Salt bridge. Salt bridge. Without the salt bridge. Is that what your question was? Okay. Without the salt bridge, the circuit won't be complete. Do you understand that? You need the salt bridge because for electrical flow, your electrical circuit has to be completed. All right? And so if you didn't have the salt bridge, nothing would happen because you have to have the circuit completed. Okay? Sorry. I miss, I heard the sulfate, not salt bridge. Okay. So we've looked at an example and this particular cell is called the galvanic cell. The galvanic cell has a particular name and this is called the danial cell. So if, so we call this the danial cell, so let me just, so if your reactants are zinc, metal and you have zinc, zinc sulfate and copper, copper sulfate, this combination of a galvanic cell has a specific name and this is called the danial cell. All right? Now you can have different types of electrodes. So let's take another example. So if you want to describe another example of a cell, let's take another electrode. So we looked at electrodes and we've looked at zinc electrode. We've looked at copper electrodes. All right? So in the danial cell, the zinc electrode forms the anode, the copper electrode forms the cathode and you have a flow of electrons. Okay? Now you can also have another electrode called the hydrogen electrode. So let's take a look at what it means when we say we have a hydrogen electrode. And let's say your hydrogen electrode is the cathode. So if your hydrogen electrode is the cathode, remember that the cathode, does oxidation take place or reduction take place? Reduction. All right? So the reduction reaction that you see at the cathode will be 2H plus equus plus 2 electrons giving you hydrogen gas. All right? And so if you wanted to set this up, your H plus comes in the solution in the form of let's say HCl equus and usually the concentration is one molar. So you can see in this, it's a strong acid, it dissociates completely. So in that solution you have H plus ions. All right? Now the product is hydrogen. So we need something. Remember we need a wire to bring in the electrons into, remember where the reduction takes place, electrons are being pulled in. So you need a flow of electrons coming into that solution. And so what we do is we have an inert metal like that and this whole, the wire as well as this is platinum and this is an inert metal. So the reason it's inert is it doesn't react. All it does is just serve to deliver the electrons into the solution and you have a test tube that looks something like this. So you have a test tube which is open at one end like that and you have hydrogen gas as well. And so essentially what happens is the hydrogen gas comes through here and then it'll bubble out on these sides. You'll see bubbles coming out and this would be the cathode. All right? So if you wanted to look at a setup where your cathode was now a hydrogen electrode, this is what it would look like. So here on this side we have the anode and the anode is the zinc electrode. The cathode is a hydrogen electrode. So essentially what you do is you have hydrogen gas at one atmosphere, you have one molar HCl, your platinum, the platinum metal is just used because it's an inert metal. You need something to deliver the flow of electrons into that solution so that this reduction occurs and so an inert material like platinum is used. All right? And you can see the bubbles of hydrogen gas. So if you wanted to describe what's going on in this cell then you would say that at the cathode, so at the anode you would have the oxidation taking place which is the zinc solid going to zinc 2 plus equus plus 2 electrons. At the cathode you would have the 2H plus ions picking up 2 electrons to give you H2 gas. And so your net ionic equation would be where you have zinc solid plus 2H plus equus giving you zinc 2 plus equus plus H2 gas. Did you guys get that? So overall the net ionic equation that takes place in this galvanic cell is shown here and you've identified what the anode is, you've identified what the cathode is and then how many electrons are being transferred? There are 2 electrons being transferred in that setup. Now you could also have the other way around. So now let's say let's take a different cell and in the other cell let's say the oxidation. So what we have is the anode is where you have the oxidation. So now let's say we have the hydrogen electrode as the anode. So the hydrogen electrode is the anode, you're going to have hydrogen giving you 2H plus equus plus 2 electrons. This would be the oxidation. At the cathode you would have copper 2 plus equus pick up the 2 electrons to give you copper solid. So this is the reduction. And now we've looked at an example where the hydrogen electrode is the anode and now the copper electrode is the cathode. So we took the Daniel cell and we took one half of it and replaced one half with the hydrogen electrode on either side. And you can see that the net reaction would be, so the net ionic equation would be where you have hydrogen gas plus copper 2 plus equus giving you 2H plus equus plus copper solid. So the overall reaction that takes place in the second example where now your hydrogen electrode is the anode where the oxidation takes place and now copper is the cathode then the overall reaction would look like that. So that's what this picture shows you. So the two examples that we looked at here is remember usually when you draw these out you're customary to always put the anode first and then the cathode second. Here you have the zinc electrode here and over here you have the hydrogen electrode. So this is the cathode and you have a flow of electrons and the circuit is completed with the salt bridge. Now over here this is the other. Now you have a copper electrode which is the cathode. Your hydrogen electrode is now the anode so now this is where the oxidation occurs. This is where electrons are being pushed out and it goes around and the electrons on this side are being pulled in. You see that? These are all examples of galvanic cells. A galvanic cell is an electrochemical device that takes advantage of a spontaneous reaction. It has to be a spontaneous reaction if you're going to have a flow of electrons to the wire. And so galvanic cell is another name for batteries. This is the underlying principle for all batteries. So now we're going to look at the last few minutes. We're going to look at what we call cell notations. So this is chemist love shorthand notations. And so we have developed this system of notations that we use to describe cells, particularly galvanic cells. So we're going to look at cell notations. And let's start with the Daniel cell. We know in the Daniel cell, the reaction that we're looking at, the net reaction, is that you have zinc solid plus copper 2 plus ions giving you zinc 2 plus equus plus copper solid. So that isn't it. And so in this reaction, we looked at the fact that the redox couples are, we have the oxidized form first, so the zinc 2 plus equus, we use a diagonal line and then zinc solid. And we have copper 2 plus equus, copper solid. That's the redox couple. All right? Now if we want to describe the cathode for this reaction, let's start with the anode because always in alphabetical order. So let's start with anode. Now when you write the anode, and this is called the notation for the electrode. So this is, we're looking at the electrode. You always write the reactant first and then the product. All right? And if there's a vertical line like that, now it's different from a diagonal line. It's very, very important that you draw vertical lines and diagonal lines and don't mix up the two. All right? If you draw a vertical line like that, it means that there, the reactants and products are in two different phases. So if you have something that's dissolved in aqueous solution and a solid, or if you have aqueous and a gas, or you have a gas and a liquid, all right? These are two different phases and what we say is that there's an interface between them. All right? So if you have a vertical line, there are two different phases. Or if you always write the reactant first and then you put a comma and then product, this means same phase. All right? So if I wanted to write the Daniel cell, my anode would be, my reactant for that is zinc. I'm going to put a vertical line because we're going from zinc to zinc 2 plus aqueous. All right? So that describes the anode. Now if I want to describe the cathode, I would write the reactant. Vertical line because again, you're going from an aqueous solution to a solid so there's a change in phase would be that. All right? Now we describe the cell, so the cell notation is always anode first, cathode last. All right? So just remember it as AC. All right? It's an alphabetical order. Just remember that. Anode always comes first. All right? And therefore, I would start with the anode. Remember with the redox couple, it's always the oxidized form, diagonal sign, reduced form. When you draw the symbols for the electrodes, the notation, it's always reactant first, product. And if there's a phase change, you put a vertical line between them. Okay? So now if I combine them to give the cell notation, what you do is you start with anode and you know the anode is that, so vertical line, zinc 2 plus aqueous. Now I put a double line and this double line refers to the salt bridge which completes the circuit. And then on the other side, we have the cathode reactant first vertical line product. So what does this symbol convey? To chemists, this symbol represents this picture. So every time you see this, that's what this picture represents. So the shorthand notation is a very convenient way to convey to everybody. That symbol, whenever you see those sets of symbol, you have to have a mental picture of what your galvanic cell looks like. And so that translates to this. So you know that the anode has a zinc electrode dipped in zinc 2 plus, alright? And that it has a salt bridge and on the other side, you have a copper dipped in copper 2 plus and it signifies this picture. Got it? Now at the beginning of next class, I will take a few more examples of writing the cell notation and show you how to write cell notations to describe what the galvanic cell looks like. Okay? So I will stop there for today.
spontaneous reaction. Remember last quarter when have an oxidation reduction reaction, you will conjugate acid-base pairs. Now in the case of which is the acid and which is the conjugate is ultimately in electrochemistry, what we're spontaneous reactions and non-spontaneous