 Hello, everyone. I welcome you to the next chapter of this chemistry crash course, which we are looking for the autonomous syllabus. The chapter's name is electrochemistry. And today, we'll look at all the different topics that we'll cover in electrochemistry, as well as how to really write the model answers. All those who are present, I request you to put in their name so that I know that you are there, as well as you can always keep on asking your doubts as such for the chapter. I'll try and answer them during the course of the chapter or maybe at the end of that. I also most welcome any comments that you might have in terms of any topic that you want me to talk on more or if you feel like understanding any portion of this, I can see that there are a few hands that I've already joined in. Please feel free to punch in your name or just put in your attendance as such. Good. So some background on this chapter before we really get into the details. So electrochemistry is one chapter where in fact, like all other chapters as well, we encounter in our daily lives almost every day. So for example, any batteries that you may use or anything that's to do with getting energy or electricity particularly from chemical energy has to deal with the concepts of this chapter. Very, very important and useful especially in these days because with the advent of solar energy and all, it's pretty critical to have this energy being stored. One of the forms in which this energy can be stored is in the chemical energy form and then it can be used or reused at any point of time. So that's a quick introduction. If you guys are around, I again request you to put in your name so that I know that you are there. So let's begin and let's understand what this chapter is all about. We'll go one by one into all the details and see how this chapter unfolds. Good. So let's get in. Now one of the most fundamental concepts to understand before we understand how electrochemistry works is the concept of oxidation reduction. As you can see on the screen, there are basically three. Just a second. Yeah. So I'm not sure are you guys there trying to really connect? I can see that the broadcast has suddenly stopped. Hello. Yeah. Why is this sharing? Let me reshare the screen one more time. Yeah. Here is the screen. There you are. OK. So once again, looking at the definitions, we know that oxidation reduction is one of the most fundamental definitions of electrochemistry. Now what does oxidation mean? Oxidation means either taking in oxygen atom or giving out hydrogen. Both of these together are called as oxidations, whereas reduction really means that you are giving out oxygen and taking in hydrogen. Now we are going to look at these definitions in much more detail in the following slides. But when oxidation and reduction are happening together, those kind of reactions is what we call as redox reactions. The very word redox comes from reduction plus oxidation. And therefore, you can say that it is nothing but redox. So this very sentence redox is from reduction and oxidation. You know that. Now electrochemistry is basically the study of interchange between chemical change and electrical work. So when both of them are happening simultaneously, whatever interactions that happen, that study is called as electrochemistry. Also what are electrochemical cells? Whatever systems we will use to have this chemical change and electrical work to be done, those systems are basically redox reactions. So and these systems actually produce electrical energy. So that's the difference between electrochemical cells. Okay, now let's go further. We know that if you have to really look into deeper, look deeper into how what are the redox concepts, then redox concepts apart from the definition of oxygen and hydrogen, we also have an electronic concept of redox reactions, which means that there is basically a transfer of electron that is being done. So redox reactions are electron transfer processes. So oxidation is when you lose one or more electrons, whereas reduction is when you gain one or more electrons. So let's take an example. For example, if we have Fe2 plus and if Fe2 plus is gaining electrons going to Fe2, we will say that this process is reduction. But on the same time, if we have Fe2 plus, actually giving out electrons to give Fe3 plus, so it has lost an electron and therefore discharged has gone more positive, this process is what we will call as reduction. The process is called as oxidation. So oxidation is when you lose one or more of electrons, whereas reduction is when you gain one or more electrons and oxidation numbers are all imaginary numbers. Basically, these are the numbers that you realize is like a pseudo charge on the molecule. Now let's try and understand balancing redox reactions. So there are multiple ways of balancing redox reactions, which we'll see just in a few slides next. Now, what are the ways to really write oxidation number? So you see any pure element, the oxidation number is always zero. For example, if you have O2 or N2 or H2, any pure element in its elemental form, the oxidation number is zero. So for all of these elements, the oxidation number for each atom is zero. Firstly, please note that oxidation number is often atom, okay? So every atom has an oxidation number. Sometimes in the same molecule, two different atoms of the same element may end up having different oxidation numbers. A classic example is, for example, if we have C double bond OH and let's say CH2 and COOH. Now, in this scenario, you'll realize that the oxygens are different in different groups. So a classic example is when they can actually end up having different oxidation numbers. In fact, particularly if I have to give you an example, whenever a peroxide is used with alcohol, so let's say on one side we have CH2OH and on the other side we have COOH, okay? So this is a peroxide molecule, CH2P. So the oxidation number of this oxygens is minus one, whereas the oxidation number of the alcohol oxygen is minus two. So both of them, you know, having being in the same molecules also ends up giving different oxidation numbers. So please be wary about, you know, oxidation numbers basically are for each atom. So when you have oxygen, oxidation number is for the atom oxygen, not for O2. For example, in SO4, two minus, which is the sulfate ion, oxidation number is for each oxygen atom, okay? So now that we know that it is for atoms, whenever the atoms are in the elemental form. So for example, when we have H2 or O2 or we have N2 for that matter, okay? So these oxidation numbers for each of these atoms is zero because they're in the elemental form. Then for monoatomic ion, oxidation number is equal to the charge. So let's look at a few examples. Let's say we had a chlorine, okay? Now if chlorine is with a minus one charge, which is an ion, you'll find that the oxidation number for chlorine is basically minus one. Similarly, the oxidation number for H plus is plus one. So and so forth. For monoatomic ions and for that matter, even in polyatomic ions, the oxidation number is equal to the charges that is present. Let's look at a polyatomic ion. For example, one example that I just gave is SO4, two minus. Now in SO4, two minus, we have O2 minus, which are four atoms. So the total charge here would be actually minus eight, whereas sulfur has a total charge of plus six and therefore plus six plus minus eight is equal to minus two. So the charge on the entire ion is equal to sum of the charges of these individual atoms and therefore that is equal to minus two. So that's the summation of oxidation numbers. Now in a neutral compound, of course the sum of oxidation numbers has to be zero and therefore oxidation number of the molecule H2O, the total oxidation number is zero. Now please note, whenever we are seeing oxidation numbers here, this is the oxidation number of the compound, whereas this oxidation number is of the ion. So oxidation number of compounds, ions, molecules, polyatomic molecules or ions, everything is just the summation of oxidation numbers of the atoms, as simple as that. When there are multiple atoms, then sum up the oxidation number of each of the atom and you'll end up getting the oxidation number of the molecule, as simple as that. Now I can see a lot many of you have joined in. It would be nice if you can just punch in your name so that I know that you guys are there. The next one is most of the times oxidation number of hydrogen is plus one, okay? And for that matter, even the polyatomic atoms, the oxidation number is equal to their charge as we have really seen. Please note only in one situation, the oxidation number of hydrogen is minus one, which is hydrides. So for example, when you have NAH or when you have LIH, so these are all hydrides here, the oxidation number of hydrogen is nothing but minus one. So that's one important difference that we have. Similarly, the oxidation number of oxygen is minus two in most of the cases, except for peroxides or whenever it is with a larger electronegative atom. So in OF2, you'll realize that fluorine is more electronegative. And therefore it should be given a negative oxidation number, which is nothing but minus one. And since there are two fluorines, the total oxidation number here is minus two and oxygen ends up getting an oxidation number of plus two. So one rule, which is, of course we have to remember is that the oxidation number generally is assigned to, negative one is assigned to the one that is more electronegative. So since fluorine is more electronegative, fluorine earns the right to have minus one as the oxidation number and oxygen ends up getting plus two. So that's one quick thing. Now, so in peroxides, oxidation number is minus one. This is minus one actually. And in OF2, the oxidation number is minus plus two. So remember these subtle differences between oxygen and hydrogen. Most of the times they would just ask you about oxygen and hydrogen, but not more than that. The other thing is fluorine also has an oxidation number of minus one. So quick thumb rule, whenever you look at any atom, I'm going to take one very complex example, maybe let's say K2CR2O7. Now you will see that there are seven plus four, almost 11 atoms in this molecule. Of course, potassium being, it shows only one oxidation state does not show any variable valency. Therefore, the oxidation state of potassium has to be plus one. Whereas the oxidation state of oxygen has to be minus two because your oxygen seems to be in its base or primary valency. And whenever we have minus two into seven, we end up getting minus 14. Plus plus one into two is plus two. So the total newt and the oxidation number of the molecule has to be zero. So the summation has to end up zero. So I have to add another 12 to it. Now, because I've added 12, 12 is between two chromium atoms. So each chromium atom, the oxidation state is nothing but plus six. So that's the oxidation number of chromium in this particular example. So that's how we find oxidation numbers. Oxidation number is pretty critical to find out what is happening in the solution, whether there is an oxidation happening or whether there is a reduction happening. So that's one important thing. Now, let's look at oxidation reduction in the form of oxidation number. So whenever an oxidation is happening, that is loss of electrons is happening, of course, you will end up getting more positive charge, which means that the oxidation number is going to increase. One example that I gave you a few minutes ago was Fe plus two. It actually giving out some electrons and being Fe plus three plus electrons. Which is nothing but an increase in the oxidation number. And since the oxidation number is increasing, we call this as oxidation. So oxidation is anything where oxidation number increases. Whereas reduction is anything where oxidation number has to decrease. So a simple concept, the same oxidation reduction, which we had seen from the examples of oxygen, we are seen from the examples of hydrogen, and we are seen from the examples of electrons, can in combination be looked at simply with the change in the oxidation number, which is going from positive to more positive or more positive to lesser one, right? So a quick rule of learning is oil rig or Leo gear, both are also okay. So again, of electrons is reduction, loss of electrons is oxidation, or oxidation, you can also remember this oxidation involves loss or reduction involves gain. So oil rig or Leo gear, whatever is comfortable for you, this is a quick acronym or something that can help you remember oxidation reduction pretty neatly. Using the same terminology with an example. So let's say, we have a simple example of zinc reacting with hydrogen, okay, H plus. So you'll realize that zinc actually ends up giving out two electrons to hydrogen and hydrogen gas is evolved in the process. So zinc turns to Z and plus two and hydrogen. So gas. Now let's look at the definitions. So oxidation means one reactant must lose electrons. If you remember oil or Leo, loss of electrons means oxidation. So one electron has to lose electrons and it is zinc who is losing electrons here. Therefore, zinc is getting oxidized. Now when zinc is losing electrons, it is giving this electrons to someone and because the hydrogen is the one who is accepting these electrons, hydrogen is getting reduced. So zinc also is the reducing agent. Because remember the reducing agent gets itself oxidized, whereas the oxidizing agent themselves get reduced. So since zinc here is giving electrons to hydrogen and hydrogen is basically getting reduced, we will call zinc as the reducing agent, okay. Now the oxidation number definitely increases and we can see that in this process, the oxidation number of zinc has increased from zero to two and it's a good, it's an increase in the oxidation number and therefore zinc is undergoing oxidation, right. So this is what it is. Now similarly, the oxidation number of hydrogen has decreased from plus one to zero, okay. So this is the change. It is plus one year and it is zero year. So you can see that the hydrogen has got reduced and the hydrogen is an oxidizing agent. Now, with this in place, just to opt for the definition, oxidation is loss of electrons and therefore it causes reduction and hence it is a reducing agent. Reduction is gain of electrons, it causes oxidation for the other person, for its counter part reactant and therefore it is an oxidizing agent. Now, these redox reactions can happen in two different forms. One form is it can happen in the same vessel, okay. Please remember the same vessel. So whenever it happens in the same vessel, we call it as direct redox. Whereas the same redox reaction can happen in different vessels, okay. So oxidation in one and reduction in one, which is when we call it as indirect redox. Now when it's happening in the same vessel, it is very difficult to extract energy out of it. So most of the times, the only thing that you can extract is heat. So heat comes out of the process and you cannot have anything else positively used. But when you have an indirect redox reaction, it is the electrical energy that you can basically use out of the system and therefore indirect redox reactions are most of the times used in industrial applications or any of our human purpose applications and electrical energy can be extracted from chemical energy, right. Now, how do we really balance these redox reactions whenever you have been given two redox reactions? How do we try and balance them? So what we do is we write what we call as half reactions for oxidation and reduction. Now each half reaction, half reaction means basically in that reaction only oxidation or reduction is happening. We're going to look at a few examples in the next slide. So for each half reaction, basically we will end up balancing elements involved in the electron transfer. So first elements are balanced and then the number of electrons are balanced. This is a very important sequence. Remember that elements has to be balanced first and then the electrons has to be balanced, okay. This is the only way to really balance them out pretty fast, okay. Now to balance electrons, we will have to multiply the half reactions. Let's say this is the oxidation half reaction and this is the reduction half reaction. We'll have to multiply the oxidation half reaction by certain entity and the reduction half reaction by certain entity so that we can cancel out the electrons. We'll do a quick example in the next slide. Yeah, once they are done, you add the half reactions to cancel out the terms of electrons which basically are not needed in the process. Now once that is done, you'll realize that once again, okay. Yeah, yeah. So now you'll cancel out. Now there could be two conditions in which oxidation reduction can happen. One is the acidic condition and second is the basic condition. In acidic condition, we know that acids means that there will be a lot of H plus. So if there is any shortage of H plus on any part, you'll simply add H plus ions to the solution and you can balance out the hydrogens and oxygens are balanced by adding H2O on the other side. Whereas in basic conditions, you'll add OH minus, you cannot have H plus in the solution and so therefore oxygens can be balanced simply by adding OH minus. Hydrogens will have to be balanced by adding H2O on the other side. Now also at the end of the equation, always check that all the atoms are basically, their charge and the number of atoms both are balanced. That's a very important thing to really know, okay. Now let's go to the next instant. Yeah, so you've seen this. Now what are the common components of the oxidation reduction, okay? The common components of oxidation reduction reactions, especially our electrodes. So for example, one second, sorry. So the common components for this is, for example, you can have electrodes. Now electrodes are all of those that conduct electricity between cell and the surroundings. Basically these are ones which carry electric current into the system and carry electrons out of the system. And then electrolyte is, it's basically a mixture. It's like a solution, which is used and which is involved in carrying the charge within the liquid. It is also something that is a mixture of a lot of ions, which is involved in the reaction mixture. And sol-bridge is something that completes the circuit basically since ions have to pass and we cannot use the same vessel which we're doing in direct applications. Therefore, we have sol-bridge which can sol-bridge which can actually pass ions from one vessel to the other. Now let's look at different types of electrodes. Of course, the two major types of electrodes are anode and cathode. Now oxidation occurs at anode, why? Because anode is the one that actually takes electrons from the substance. So if you are an anode and let's say you are taken from our previous example, we have Fe2 plus. So as you reach the anode, you'll realize that anode is the one that can take electrons from you if you are in the electrolysis, electrolytic cell and therefore oxidation occurs. So anode takes electrons and therefore oxidation always occurs at anode. And cathode actually gives electrons. Cathode is the one which is negatively charged. It gives electrodes and therefore it is the reduction that happens at cathode. Now there are two different types of electrodes also. One is called as active electrode and then active electrodes. Active electrodes are the ones which participate in the redox reaction. For example, you use a copper electrode or you use a zinc electrode. You'll find that copper will dissolve and, sorry, zinc will dissolve and it will get electroplated on copper, okay? So copper gets plated on it. So basically though electrodes themselves are participating in the reaction, they are called as active electrodes. But let's say if you take carbon electrodes which has to do nothing with the reaction, they simply are passing current. Then these are called as inactive electrodes. They are simply sites of oxidation reduction, but they are not the ones who themselves are taking part in the reaction. So there are two different types of electrodes depending on whether they are a part of the reaction or not, which is active electrodes and inactive electrodes. Now, like we have seen types of electrodes, there are also two different types of cells. One is what we call as the voltaic cell or the galvanic cell. And second is what we call as the electrolytic cell. We'll see electrolytic cell in after a few slides. Right now we are just looking at the voltaic cell or galvanic cell. Now, what is a galvanic cell? This is the device which converts chemical energy to electrical. So whenever chemical energy is converted to electrical energy, you'll realize that this is basically your galvanic cell. And the reverse one is when electrical energy is converted to chemical energy, it is what is your electrolytic cell. Okay, so now most of the times you'll find that your batteries are nothing but galvanic cells. Why? Because the chemical energy that's stored inside the battery is basically used to power up your remotes or your watch or any other electrical device. So that's your voltaic cell. Now, in voltaic cell, it is very important to have spontaneous reaction. Now, let's understand what do we mean by spontaneous reaction? You have seen a glimpse of this in thermodynamics. Spontaneous reaction are those reactions which happen by themselves. You don't have to really give them any energy or stimulate them or have to use any external force to make the reaction happen. So these kinds of reactions are called as spontaneous reactions. These reactions are used in galvanic cells to give you energy. So that's the basic advantage of having spontaneous reactions. Now, here's a quick example. I actually wanted to show you some figures. This is one of the advantages of an online program is that I can actually show you. You see, this is a zinc metal wire. Now, this zinc metal wire basically is dipped inside a copper sulfate solution. You can see that these are the atoms of zinc. Zinc atoms are basically getting dissolved. Why? Because they want to give electrons. And as they give electrons, they turn up into a Zn2 plus cation. And therefore they are the ones who are getting into the solution. But as they are getting into the solution, they give these electrons to the incoming copper. Now, this copper was a copper cation earlier which turns into a copper atom. And since it turns into a copper atom, eventually it starts getting deposited on this zinc ions. So you can see that this is how the copper has got deposited. So zinc is giving electrons to copper and copper is getting into the metallic state. Please note the states. Zinc is in the solid state earlier. Now, it is in the aqueous state. Aqueous state means it has dissolved. Whereas copper was in the aqueous state, now it has become solid. Copper has basically got deposited on the zinc plate. So this is a quick example of how it really happened inside the solution. Now zinc, but there is something that we also have to understand which we'll see in a few slides again is that, why is it that zinc got dissolved and copper does not? Why is that zinc is going into the solution? Copper is coming out of the solution. This actually depends on reactivity which we'll see in a few slides to come back. Now, whenever such spontaneous reactions happen, now here you see both zinc and copper are in the same solution, which means this is a direct redox reaction. I cannot really do anything about it. The zinc is getting dissolved, copper is getting deposited, or everything is happening by themselves. It is no use for me because I will not be, excuse me, I will not be able to extract any energy out of this, whether electrical even the heat energy is pretty nominal. So that's why we use what we call as the Voltaic or Galvanic cell. So this is one of the setups. In fact, here's a quick image. So you see, this is the copper which has got deposited on the copper rod, whereas this is the zinc which is getting dissolved. So what we have done now in the indirect redox, which is different from the earlier one, we have kept separately zinc and, zinc as well as the other copper electrode and their beakers are also separate. So both of them are completely different from each other. Now, zinc is getting dissolved as Zn plus two, but it is giving electrons. These electrons are not being given to copper directly. Please see that these electrons are given through an external wire to copper, whereas copper here is taking those electrons and getting deposited onto the copper strip. Now, because copper is getting out of the solution, it needs more positive ions in the solution because all the positive ions are actually getting deposited as copper here, whereas here there are a lot of positive ions getting formed. So since more positive ions are getting formed here, there are more negative ions that are needed here and more positive ions needed here. You can see that the sodium from the solid bridge is actually coming to the rescue of all the negative ions and balancing the charge out, whereas SO4 2 minus is coming to the rescue of zinc so that the solution remains electrically neutral. We cannot keep the solution getting charged because after a certain amount of time, a reverse potential will get generated and we will not get any potential out as the voltmeter. Okay, so now in this indirect reaction, since in this cell you'll see that only oxidation is happening, which means only zinc is going to Zn plus 2, this is called as the oxidation half cell, whereas in this beaker, there is only reduction happening because CO is going to copper, activist form is going to the solid form. So this is the reduction half cell. So this is how in indirect redox reactions, the oxidation half cell and the reduction half cell can be separated out and both of them can actually be individually assessed as well as we can understand what's really happening in each of these beaker to extract the right amount of energy. Now, since these electrons are moving from an external circuit, these electrons can be used to do some work. This is one of the biggest advantage of this indirect redox reaction. Now, if you really add these half cells, you'll realize that you end up getting an overall cell reaction. Okay, so this is one of the cell reactions that can be found pretty commonly used in fact, this cell is also called as a Daniel cell and it is commonly used to experiment or to study electrochemistry. It is, please note that this is also a Voltaic cell or Galvanic cell or an electrochemical cell because this is chemical energy getting converted into electrical energy. Also, it is an indirect redox reaction and oxidation cell and reduction cells are different. Good, now, here's a quick look at how the anodes and cathodes had really looked out in this. So you'll realize that zinc, this is basically the copper one now. Please note that this was the copper electrode. Now, in the copper electrode, the reducing agent, so there's an oxidation that was happening in this. Sorry, this is the, yeah. So if you can see that this is the cell where there was an oxidation that was happening. So this is the zinc one. Okay, I'm sorry, this is the zinc one. So in the zinc electrode, you'll find that oxidation was happening because zinc was oxidizing. These electrons were going to copper and reducing it. Therefore, this zinc was a reducing agent inside the oxidation half cell, whereas since the electrons went here and it was the copper two plus, which actually accepted that there was reduction happening in this partial half cell. But because it was taking an electrons, it was effectively getting zinc oxidized. So it is an oxidizing agent. Of course, you know that reduction and oxidizing agent or oxidation and reducing agent are in the same vessels. Now, because of this reason also, because the electrons move out of this and they are coming in here, and these are the anions that, since electrons are moving here, this is the anode and this is the cathode. We'll also look at the charge in a short bit, how to really mention the charges. Now, what are the basic concepts? Anode is the one that is from where electrons are actually, so if you really have external bulb or something out here, since the electrons are coming from here, you will realize and since the electrons are going in here, let's say if this would have to be replaced by a battery, you'll realize that the battery would have been something like this. So electrons would come in from the negative side and electrons would go into the positive side. So the cathode is the one which where electrons are coming out. So this sense it is called as cathode. So the very definition of cathode and anode is only based on the solution definition. This is very important, please note. Cathodes and anodes are based on the solution. Please do not get confused with what is happening outside the solution. Outside of the solution, you will find that the electrodes are completely different. The charges are completely different. I'm just gonna show that in a quick minute. Now, when zinc and zinc metal were kept in a direct, you know, a redox reaction, the zinc plate simply disappeared, you know, and the copper were simply getting deposited. So zinc strip disappears, whereas the copper simply keeps on depositing. Therefore, Zn was getting oxidized and it was the reducing agent. We've seen this in the previous case as well. Now I'm trying to write down the equation for it. Also, the other case, if I have to write the copper equation, copper is actually getting reduced and you know, it is the oxidizing agent. So these are the reactions, half reactions that have happened. You have seen the same reactions in the previous case as well in the oxidation half cell and the reduction half cell. Yeah, so now let's look at, you know, how to really obtain a full current. So to obtain this full current, we had taken a direct redox reaction into an indirect one with the sol-bridge, basically. Now we look at what really has happened in all of these processes. So firstly, we achieved electrical current through a galvanic cell, which is also a part of battery. So what has happened through this entire process? Firstly, as zinc gave Zn2 plus, this is the observation. So if you have to set up this whole tech cell, the observations are, zinc gives out Zn2 plus. You have to write all of these in your answers. It's a good idea to actually draw the label diagram. And of course, as you go forward, you'll realize that, sorry, on the other side, yeah, as zinc actually gives electrons. On the other side, copper takes in these electrons to form copper metals, right? Now, the oxidation, it happens at anode, but please note the charge. The charge is negative. Why? Because as I mentioned to you, since electrons are going out of this once again, yeah, since electrons are going out of this, and let's say you had a bulb here once. Yeah, if you had a bulb here, right? The electrons are coming from this bulb. So you'll realize that the circuit for this scenario would be that as if you had put in a battery which had its negative end at this end and positive end at this end. Therefore, although this is an anode, why anode? Because the electrons are actually going to this, right? So electrons are going and Zn plus two is coming out. So as electrons are going, it is an anode. Please note that the definitions of anode and cathode as we are discussed is only because of what is in the solution. And since for the external circuit, it has a negative charge. So this is the ambiguity and this is very different from electrolytic cells. I had mentioned this in the class as well. Anodes are negatively charged whereas cathodes are positively charged. Let me show you the cathode expression as well. Yeah, so reduction is happening at cathode and it is positively charged. So remember for electrochemical cells, red cat, okay? Red cat is an acronym that you can remember which means that reduction happens at cathode and it is actually positively charged, okay? Now, what are the different uses of salt bridge? Okay, so there are basically four or five uses. The first one is it is responsible for movement of ions. Okay, so first one is simply responsible for movement of ions. Now, why does we have ions move? These ions move because they want to balance the charge or maintain the neutrality. Maintain the neutrality of the respective cells. What is the third one? It actually wants to kill the reverse potential, okay? Or destroy the reverse potential. What do you mean by destroy the reverse potential? See, if these ions do not move, there will be too much of positive charges here and too much of negative. Right, so I hope you guys are able to see this. Hello, yeah, coming back. So yes, so we were talking about salt bridge. So we realized that salt bridges gives us multiple benefits, it actually only have, yeah. Sorry for the minor sound. I know that one issue. Yeah, so the first one is of course it maintains, you know, it makes the ions move. The second one is coming back and joining. I guess now you can, yeah. So sharing you the screen one more time. Yeah, so getting back with our concepts is that we still have, we're looking at the salt bridge. So salt bridge helps in movement of ions and it also helps in keeping the neutrality of the molecule. Yes, now, let's go further. Now, what are the other terms used in the voltaic cell? A few of the terms is, you know, so this is an entire layout of how a voltaic cell really works in terms of the schematic diagram. So you'll realize that the electron flow direction is this way, but please note that the current flow direction is this way. This is the positive electrode, which is nothing but the cathode, as I've been saying all the time again and again. This is very named cathode comes from the solution because, you know, in the solution, the electrons are coming from this electrode, so therefore it is cathode. Whereas in the solution, the electrons go with this electrode and therefore it is anode. And please note that it is negatively charged. The, you know, the reduced species come and they actually give electrons out here and they become the oxidized species. So here is where oxidation is happening. As we have mentioned here at the bottom, this is the anode compartment or also called as the oxidation half cell reaction. This is the cathode compartment, also called as the reduction half cell. So both of them are happening at the same time to have the reduction reaction. The electrolytes in the, is the ions in the solution on both the sides. Now, please note in the salt bridge, it is the ions which is moving through the salt bridge. So here we have written it in a very short form, although it is an entire bridge that you can draw for it. But without drawing the bridge, you can just simply mention this way. So now here we will understand how do we really write the cell reaction for this or was representation for this. The cell representation for this is actually done through writing from the oxidation to the reduction one. Always remember OR or OR, that is always OR how you write the cell representation. So this cell is represented as, you know, zinc going to ZN2 plus. So whatever is happening in the solution, you have to write in the same form. What happened to zinc, zinc went from ZN2 plus. So I'm writing it in this form itself. Then you actually write two double bars and then you write CU2 plus was going to CU, okay? Now having said this, please note that there are two more very important things to write. One is the state, okay? So in what state or what phase it is. So always write this aqueous, solidier. You will write this as aqueous and another solidier. And the second thing that you have to write is the molarity or the pressure in case of gases. If it is simply atoms or ions, you write the molarity. So for example, here we can write this as one molar. Here you can write it as one molar. Had it been gases, then I would have to write the pressure. So this is a way that the cell representation happens. Cell representation always happens from oxidation to the reduction side. And they are basically distributed or cut by two double lines. These two double lines is nothing but the salt bridge. Sometimes if there is a salt bridge present, you can write that as KCl, agar-agar or whatever it is inside the salt bridge. So that's a quick look at the cell representation. Now, what really means when galvanic cell is working? Let's look at this point in slightly more detail. So you will realize that Zn2 plus is going to Zn, which means that it is trying to give up electrons to copper. So it's actually putting in energy effort. It is trying to push harder on electrons to be given to copper. Why? Because copper also wants to give out electrons. So who will give electrons is basically decided upon who is stronger in giving electrons. Now, because Zn has more potential energy to give electrons, it wins. And therefore Zn ends up giving electrons to copper just because of the greater potential energy. And therefore we also call it as greater potential, nothing but greater potential. On the other side, in spontaneous reactions, copper loses because copper has a lesser potential. Now, why does this spontaneous reaction happen because of the relative difference in metals' abilities to give electrons and the ability to make the electron flow through the solution? So someone who is very eager to give electrons will want the electrons to be given to the next atom. And therefore, spontaneity comes in play. Please remember, spontaneity has nothing to do with time. It is not something that we are looking at time. It is only the thing that we are looking at in terms of the energy that it has. Can it really give energy from one to the other? If that is what it can do, then it will really work. So it has not to do with time, it is simply to do with the energy. Remember this very one important point of spontaneity. Now, let's go further and understand what is the cell potential. This very cell potential, where we were speaking in terms of... So I can see a lot of... Your messages have just popped in. They were not visible earlier. So I can see all of you. Hi, Neha, Brian, Kirtna, Ronak, Ananya, Aditi and Sanchi. Good to see you all, guys. Now, if there is anything that you're not connecting on, please feel free to message. Now, cell potential is basically the drive or the force. It's actually an internal energy that the element has to give these electrons. And it can actually be measured through a potential difference with any of these electrodes. Now, we'll see how to really measure these difference. So it can be measured by what is the electrical potential energy or a work that the cell does in moving one charge. So this is how voltage is given out. Voltage is nothing but concentration of energy. We have seen this in the earlier classes as well. So what is the work or concentration of electrical energy per unit charge? That's what is electrical or cell potential. Now, in our situation, let's say zinc, it has a... For one molar solution of zinc, if it is giving electrons to copper, you get a reading of 1.110. It means that zinc is 1.110 volts more than copper. Why? Because it's going to oppose taking those electrons also. Copper also wants to be in the solution phase. Please remember, all the ions always, as much as possible, would like to be in the solution phase. Because of some very, very less reactive, for example, ion or gold, they are so less reactive that even at room temperature, even at normal conditions, they end up forming solids. But given a chance, every element, every ion, would want to be in the solution phase. Why? Because the degree of freedom or the energy is very high in the solution phase. That's why they actually dissolve. Now, having said this, how do we really measure the cell potentials? So the cell potentials are measured under standard conditions, which means that the temperature has to be 298 Kelvin. The solution has to be one molar. The atmospheric pressure has to be one atmosphere when it is in terms of gases. So this is a way that cell potential is measured. Please note that there is nothing called as electrode potential. Electro potential is a very relative term. It is basically the potential of a cell with respect to someone whom we have set as zero, which we'll shortly look at. So for the time being, what we are saying is that your cell potential is the potential of the cell, which means it needs to have oxidation and reduction, both have cells at these given conditions. And we measure it with a voltmeter and we'll get some reading on it. Now, whenever we have in this cell one electrode as hydrogen electrode, then we say that whatever potential is coming up, that is the potential of the second electrode. Okay, as simple as that. So if we have zinc connected to hydrogen, then whatever potential we'll see on the voltmeter will be the potential of zinc only. Why? Because we have arbitrarily, without not much of thought. I mean, what I'm saying is it's like a convention, without really, like a relative scale, you start somewhere as zero. So we said that the hydrogen electrode is something that will always have zero potential. Now, when will it have a zero potential, the conditions are written on the screen. There are a few very important parts. One is that the atmospheric pressure of hydrogen has to be one atmosphere when it is being blown into the hydrogen electrode. How is the electrode made? It is basically a glass tube where there is a platinum inside it. Why platinum? Because platinum is a very inert electrode. And therefore platinum is put inside this. As there is platinum put inside this, there is a platinum wire that we have. And this platinum wire is the one which is going to exchange electrons with hydrogen in the solution. You use one molar H plus, which means there is some acid that we are putting. You see, how can only H plus be present nodes? So we cannot, right? So we need to have HCl or H2SO4. The idea is with HCl and H2SO4, you also end up getting negative ions. So therefore we want to avoid these negative ions. So in a very hypothetical ideal situation, we will use something like a very weak acid, but making sure that H plus ion concentration in the solution is one molar. So the negative ion should not really bother us in our experimentation. At the same time, the H plus ions should be about one mole per liter. At that concentration, when we connect that electrode to any other electrode, electrons are getting exchanged between one another. And as they are getting exchanged, you will realize that whatever reading comes up here is called as the electrode potential, okay? Which is the, you know, now please note that this can be a positive one, you know, in the same, I mean, of course, voltmeter cannot really have a negative potential as such. But there could be situations where hydrogen is getting reduced. For example, in this scenario, hydrogen is taking in electrons, but there will be situations where hydrogen is actually giving electrons. So you will end up, you know, this oxidation reduction cell as the SHG, the standard hydrogen electrode, as it is represented in our textbook as normal hydrogen electrode is the one that is getting oxidized is the oxidation of cell and the other element or other electrode is getting reduced. In this scenario, we'll call that potential as negative. So please note that the reduction potential of hydrogen is always used as zero, not the oxidation, which means that the hydrogen has to be on the reduction side all the time to have a positive potential. So in this scenario, since zinc is on the oxidation side, we will say that the electrode potential of zinc is 0.76 volts because zinc is getting oxidized and hydrogen is getting reduced. I hope this point is very clear. This is one of the most important points people make mistake here. When the oxidation reduction cells are exchanged and this hydrogen becomes an oxidizing cell, then any potential reading that you will get in this voltmeter is called as a negative electrode potential. Why? Because hydrogen is actually getting oxidized in that scenario, right? So how do you find standard electrode potentials of electrodes? Now, see, there is no potential of electrode. As I mentioned, it is a relative potential as with respect to hydrogen. So what we have is that, you know, these are E naught values or standard electrode potentials and these are at one molar concentration and one atmospheric pressure, you know, measured with hydrogen. So you'll see if you measure copper, copper gives out, you know, four 0.34 volts with, you know, standard hydrogen potential. But see, look at copper. Copper is actually getting oxidized and therefore the electrode potential of copper is going to be plus 0.34 volts. Now, if you really look at, you know, SO4, 2 minus and 4 H2E, you'll see that hydrogen from H plus, it remains hydrogen. But here, you know, there is no change in oxidation number, but here the hydrogen has got changed in the, you know, in both of these scenarios, it is still plus one, plus one, okay? But let's look at the sulfur, okay? In sulfur, it has gone from plus six, which is here, to being a plus four here, okay? So this is six plus two. So it is a plus four. So sulfur has got reduced. Because sulfur has got reduced, hydrogen has got, okay, let me quickly see. Oxygen, yeah, oxygen has got oxidized, actually, okay? Yeah, so since sulfur has got reduced, the electrode potential of this actually has to be considered negative, okay? So this E naught will be minus 20. So please remember that so long that, you know, your, the counterpart is getting oxidized, the oxidation potential is positive. If it is getting reduced, here sulfur is getting reduced. So the electrode potential of sulfur has to be negative, minus 0.20, okay? I've just mentioned the values here. The sign is not yet mentioned, okay? Now, let's look at a few values as such. Now, if we really look at, one quick second, yeah. If you look at the cell potential E, for ZNCU cell, the potential is 1.1 volts at 25 degree Celsius, when zinc and copper was at one molar concentration. This is the standard electrode potential of the cell. A quantitative measure of this tendency is given by this electrode potential, which means how eager is zinc with respect to copper to really make reactions happen, okay? So how stronger is zinc with respect to copper? That's what this electrode potential really is mentioning, okay? So it's a quantitative measure of the tendency of reactants to produce products in this scenario. Okay, now let's go to next. Yeah, now, how do you calculate the cell voltage? Firstly, you balance the half cell reactions, which are, can be, they can be added together to get a balanced equation. So in our scenario, this is the half cell reactions that we have. So since copper is going to see you solid, you simply add up the half cell reactions and you'll end up getting the complete full cell reaction, okay? So this is how we calculate cell potentials. Now, when you are doing and adding this, you are also adding their, one second, you are also adding their potentials, okay? So for each of these, we know the E naught half reaction. So we can simply get the E naught for net reaction. Let's look at some examples. Here, the E naught is plus three, four. Hydrogen is zero. Zinc, zinc, standard oxidation potential is positive. So reduction potential is going to be negative. So you'll realize that copper has a better oxidizing, it is a very good oxidizing agent because it is oxidizing zinc, but zinc is a very good reducing agent, okay? Now, if you go further, you'll realize that the reducing ability of the element is more for zinc, as we just mentioned. Now, if you have to find the cell potentials in the electrochemical cell, we can simply take their value. So you see that at anode, since it's the negative electrode and it's the source of electrons, you'll find that zinc's electrode potential, we have written positive, copper's electrode. Now, please note, these are all reduction potentials, okay? These are reduction potentials. Now, as I add them up, you'll realize that I will have to add their potentials themselves. If I reverse the reaction, I will have to change the sign. So the oxidation potential is simply the, you know, the negative of the reduction potential. So as you reverse the reaction, you simply change something from positive to negative or negative to positive. So finding the electrochemical voltage or the electrode potential potential difference between both electrodes, you are simply going to add the reactions in the right form. And as you add them in the right form, you'll realize that, you know, their electrode potentials also are added. Now, please note, if you multiply this by two, your potentials should not be multiplied. That's a very important thing. Let's say in some places you have two C U and four electrons or two C, your electrode potentials should not be multiplied, okay? Because electrode potentials are not what we call as an extensive property. It does not depend on the amount of substance. It is something to do with concentration. It's like concentration, right? So you can have four liters and two moles, which will also give you 0.5 molar solution. And you can have 0.5 moles in one liter, which is also 0.5 molar solution. So it's like concentration. Concentration does not depend on amount of substance, but it depends on the amount and volume both. Now, here are a few, you know, major electrode potentials that I just wanted to show you of different reactions. So you can see that fluorine taking electrons actually give 2.87. Similarly, oxygen with hydrogen, you know, forming oxygen gas plus water, the, you know, the electrode, this is basically the E cell, okay? So the E cell of the reaction is plus 2.07, which is pretty high. You can see that the, now, obviously, you know, if you're using a very low reactive element and a very high reactive element, as far as these elements are further apart in the electrochemical series, you are going to have a very high, you know, cell potential. Now, let's look at some very odd ones. If you really see HNO3, so this is NO3 minus and for hydrogen, when hydrogens actually connect with NO3, you end up getting NO plus H2O, which is nothing but 0.96, you know. So this is one way of, I mean, this is another cell reaction and the potentials are mentioned. Just for a quick understanding of the cell reactions, yeah, please note that the hydrogens, you know, potential, we always take it as zero, okay? So the potentials that are above hydrogen are positive, potentials that are below hydrogen are negative. The potentials of hydrogen are only zero, okay? So the other potentials are not, okay? Now, so this is a quick, you know, electro potentials map that you can actually have. Now, let's look at the voltages of some voltaic cells commonly used around our houses. So if you use the voltaic cell, like a common alkaline battery, you will find that its voltage is about 1.5 volts for a lead acid car battery that is used in the cars. The voltage is about two volts, you know. So six cells generally are about 12 volts, okay? Then, you know, the battery calculator that you use is somewhere around 1.3 volts. Or these are basically cells using all of these voltages, right? Now the yield fish that we have, electric yield, it generates somewhere around 0.5 volts. Then, you know, if you have a giant squid, it generates about 0.07 volts. So these are generally some voltages that I just thought might interest you and therefore I thought of really sharing these some voltages with you. Now, we have seen so far everything about electrochemical cell, okay? Or what we call as galvanic or voltaic cell. Now let's look at what we call as electrolysis, okay? So in galvanic or voltaic cell, chemical energy is converted into electrical energy. Whereas in electrolysis, electrical energy gets converted into chemical energy. So whenever you are forcing a current through a cell to produce a chemical change, you know, you will find that the cell potential basically is negative. Which means, see, if you take two elements and you are getting your cell potential as negative, it implies that the spontaneous reaction is not going to happen. It implies you are going to be needed to supply the energy to the cells so that they actually work up, okay? So forcing a current through a cell to produce a chemical change will always end up having a negative cell potential. Now, please note this is a very important fundamental understanding that whenever you take two electrodes or two half cells, you will always end up getting some potential. There has to be some potential coming out. Only in certain situations, you know, this potential has to be very low, you know? But having said that, there will be only one way that the reaction will happen. Now, when you want the reaction to happen your way, which is the opposite of the way that it is happening, that is when you call it as electrolysis. So, you know, if I have to really write this down here, so, yeah, yeah, yeah, yeah. You know, I can just show you an example here. So you see, I have your, you know, anti-money. Sorry, yeah. So I have tin and copper, okay? So you'll realize that in tin, since it is, you know, by naturally, it is one which is going to donate electrons much more than copper, you will find that copper should be taking in electrons. So naturally, this reaction proceeds from tin to copper and the electrons move from tin to copper through external circuit and through internal circuit from copper to tin. Now, having said this, if I don't want the reaction to go in this format, I will put an external source greater than 0.4. Please note, till I'm not putting it till as much as 0.4, whatever I put, let's say tomorrow I put 0.40. It implies that only a cell voltage of 0.08 volts will happen and the electrons will move from tin towards copper, okay? But at the same time, it will still keep on opposing. If I put at exactly 0.048, I am just going to stop the reaction, okay? And when I'm putting more than 0.48, I am using the external electrical energy to actually convert or get stored into the chemical form. So this is called as the electrolytic cell, okay? Where external energy is used to really store energy into the chemical form. And voltaic cell is when chemical energy is used for electrical forms. Now, forget the delta G and E cell values, this is just from a thermodynamic perspective, but some very important differences between voltaic and electrochemical cells. In voltaic cells, please note that the oxidation process, so oxidation reduction always happened at anode and cathode, that's not a problem. But in a voltaic cell, anode is negatively charged and cathode is positively charged. Whereas an electrolytic cell, typically as we do for any reactions, anode is positively charged and cathode is negatively charged, right? So remember this difference that you have between anode and cathode and their signs between voltaic and electrolytic cells, okay? Now, let's look at the electrolysis of some typical compounds. For example, here we are looking at the electrolysis of water. It's a very interesting example. In fact, it's good if you can actually see a video or two of this on YouTube and you'll find that it looks pretty very interesting, very nice. So you'll find that in a normal electrolysis, you can actually use the downward displacement of water to collect hydrogen and oxygen. So you simply put electrical energy inside the solution and hydrogen is collected over water and oxygen is collected over the another thing. Now, hydrogen is always, if you see hydrogen is at the oxidation, half reaction and as we know that oxidation always happens at anode. So this is your anode and at the other end, it is your cathode, you know? So the reduction half cell, it is the reduction that is happening at cathode and you'll realize that oxygen will be giving out electrodes, okay? So, sorry, it's the oxygen that is getting oxidized and therefore oxygen will always be present at anode and hydrogen will be present at cathode, taking up electrons from it. And the overall cell reaction is actually going to twice of hydrogen plus oxygen. The same electrolysis can also happen with aqueous KBR, okay? With aqueous KBR, you'll find that, you know? So this is a solution of potassium bromide. You can see that this is bromine here, you know? And as you are electrolyzing, you know, you'll find that bromine starts getting deposited of KBR and potassium comes and gets deposited on the other side. So this is a quick picture of how bromine and potassium get separated as you do electrolysis of aqueous KBR. Now, you know, your alumina thermite process or extraction of aluminum, that's also something that you will have to know. This is one quick slide that I thought I would also involve because they're involved in electrochemical process. So in the electrochemical process, you will find that, you know, aluminum is basically getting, so this is, you know, aluminate, in fact, AL2O3, which is bauxite, you know, fine with cryolite. If you remember, cryolite was used in the process. So also please check how the design of the electrolytic cell is being done, okay? In this design, you will find that the lower end of the tank is slightly curved so that, you know, it actually flows out. And as it flows out, you know, at a plug it can actually be collected. So that's a way. And molten AL2O3 or Na3AlF6, this mixture is getting electrolyzed. You can also see both the electrodes at the top and the bottom one. And there is carbon dioxide that is getting formed at the anodes, whereas at the carbon-lined ion tank, there is aluminum that is getting deposited. Now, just for your understanding, you know, I've not really written cathode and anodes here. If I may have to ask you what were cathodes and anode, we can actually write them as follows. So since aluminum is coming out here, obviously this has to be cathode, okay? And this has to be your anodes, okay? Graphite rods are your anodes. And you'll realize that anodes, you know, CO2 is being given out. One second, it's not visible, yeah. And at cathode, all the aluminum is getting deposited, okay? So this is, the cell wall is lined. So this is your cathode, basically. And this is all your anodes, right? So that's a quick way to really look at the aluminum, you know, aluminum thermite process or simply sodium aluminate. Now, all of these reactions what we saw, you know, are part of a system of cells which is called as battery. So battery is nothing but a galvanic cell or more commonly, it's a group of galvanic cells connected in series, please note in series, why? Because each one has to contribute to the next one to really get the output as we need it. So these are called as batteries. Now, the next is to show you different types of batteries. So this is a typical dry cell, you know, it has a positive terminal and a negative terminal. One second maybe, I can just, yeah. And, you know, there is an insulator. Inside the battery, if you really see there are graphite rods, you know, so these graphite rods is something so that they do not degrade or they do not really dissolve, right? And then there is MNO2 and H4ClZ and CL2 paste inside it. And then there is a zinc anode. So please note that, you know, the outside of the battery is also something that is pretty much charged, okay? So dry cells are these, which where, you know, the electrode is inside and the anode is outside. They both are connected to your apparatus, whatever you are using. This is an alkaline battery, you know, which also we use as cells, you know, we call them as pencil cells or small cells, mini cells, whatever you can say. So there is a zinc and MNO2 as electrolyte in a KOH paste. Then there is a steel casing that we use. There is an absorbent. The negative end is at the bottom, positive end is at the top. Please note that there is a gap here, okay? Why the gap? Because, you know, at the positive end, electrons, you know, should get, should start coming out. And at the negative end, electrons should, you know, start going in, right? So the electrons are coming from the negative end. So the electron should get to the negative end and they should come out of the positive end. So, yeah, so that's a quick look at alkaline battery. These are just some examples so that you really have a connect to them, nothing more. This is a lead acyl battery, you know, we have alternate plates of positive and negative electrodes. The positive ones are cathodes, which are lead grills and the negative ones are, you know, anodes which are, you know, spongy lead. So positive ones are PBO2 and negative electrodes are, you know, simple leads, you know, spongy leads. Okay, good. Now the last topic of this chapter is actually corrosion. So some metals such as copper, gold, silver and platinum are relatively pretty difficult to oxidize. Okay, they are also called as noble metals. So corrosion is that property where metals actually get oxidized in presence of moisture or oxygen and, you know, they end up forming their oxides or hydroxides. So that's when, you know, we have to protect them from vapors, water or simply oxygen for that matter. So whenever you have metals which get oxidized, there are multiple ways that you can protect it. Copper, gold, silver and platinum are difficult to protect therefore they are called as noble metals. That's one thing. Now let's look at a few more ways. So corrosion of iron, okay? So, you know, I'm just taking this opportunity out to, you know, really show you how corrosion in iron happens, you know? So let's say, you know, this is the iron metal, okay? So I've just shown a little iron metal here. So this iron metal actually gets dissolved into the water as Fe2 plus while it is giving electrons. The electrons that it gives is taken by the hydrogen because in the water there are a lot of H plus ions, H plus and OH minus, you know this. So the H plus ions are the ones which are taking in electrons and you will find that this actually ends up becoming a cathodic region. So to your surprise, you know, rusting actually is like an electrolytic cell and as we have this electrolytic cell, you know, in very small quantities. So there was a very interesting thought that someone in one of the classes represented that can be actually extract energy out of this electrolytic cell, you know, in a very small form, although it is very little, you know, it still has some energy. So basically what is happening is that the iron gets dissolved and it is giving out electrons which is taken away with H plus and these H plus then basically, you know, turn to forming H2O molecules. It is the oxygen who is taking these two electrons and forming water and the Fe2 plus actually migrates through the drop and reacts with any remaining oxygen and it ends up forming Fe2O3 in the process. Now, remember oxygen forms, you know, a small pit in there and it actually keeps on accumulating more and more of this hydrogen. So the nature itself helps out in the process, you know, so pit means, you know, this is how the small, you know, because as more and more iron dissolves, you'll find that there'll be a small pit or a small hole that will start getting created and more and more water can then be stored and as more water can be stored, the process starts becoming more and more efficient. So therefore, this is a very typical way of how rusting happens and rusting is a more efficient process. In fact, it is an exponential process. Earlier the rusting would be slower, but as small iron rusts, it becomes faster and faster with time. So that's another interesting thing about rusting. Now, how does a metal metal contact happen in the rusting of iron? This is basically in the perspective of actually what we call the sacrificial coating. So zinc coats, you know, the lower metal sacrificially. You'll see that copper, if we, you know, coat it over iron, it is the iron that will still keep on getting rusted. Why? Because copper is very unreactive, very inert and therefore it does not really function as any cathode anode. At the same time, if you put zinc, zinc starts becoming the anode instead of iron. Why? Because zinc is more reactive. So zinc starts making all of those movements inside the water droplet. And there is a small potential that is created inside the water droplet. And zinc starts getting to ZNO. The good part about the ZNO is that it actually ends up forming a layer over the iron. And once it forms layer, no more, you know, zinc also can not react anymore because the layer avoids the contact of zinc with water or oxygen. And hence, you know, the corrosion stops. So these are two examples that I've shown in terms of enhanced corrosion and cathodic protection. So what we call as a sacrificial corrosion is basically mentioned technically as cathodic protection. And if you actually cover with copper, you will find that since copper is avoiding the surface area, water will spread more towards the iron and it will corrode much faster. Therefore, if you have a copper strip connected to iron, you will find that the corrosion is much more faster. It enhances the corrosion. So that's another interesting aspect. We only know about sacrificial stoppage of corrosion, but there is also something that can increase corrosion, something like copper. So these are some very interesting metal-metal contact and corrosion of iron methods. So this actually brings us to the end of this session. But I'm going to pause here for a minute or two and I just want to ask all of you if there is anything that you would like to do. This is a time that we can do. If you want to solve a problem or a two or you have any questions to ask, I'll be very happy to answer them. I'm just going to pause here for a minute and wait for your comments. Please feel free to push in your comments and ask up. I will try and help anything that you have a doubt on, any questions, problems, post which I need. I just need five minutes to really talk on the important aspects of this chapter and how can you write answers to those. Yeah, please feel free to post in your comments so that I'm aware about the same. Yeah, hi guys, I'm still waiting up. So if you have anything to ask up, I think I'll be or else, okay. Good, okay, so there are some doubts that are coming up. Let me just, yeah. So just hang on guys, there are a few questions that I can actually answer. Okay, so I'm going to talk about a few points here. Common mistakes that students make because I think people are still writing on the doubts, but let me at least cover the points that I had in mind. So please note some mistakes that I have found. If you have a reaction, for example, let me talk about, okay. So let's say calcium CaCl2, or rather I would say that I have a reaction where Ca ends up being Ca2 plus plus two electrons and I have another reaction where I have, let's say, yeah, nickel, let me take nickel. So nickel plus, plus electron ends up being nickel, okay. Now let's say the electrode potential for this thing is 0.4 and the electrode potential for this thing is 0.6. Now, obviously you will have to multiply this with two before you add them up to give you the net cell reaction, which is nothing but Ca plus twice Ni2 plus, giving up Ca2 plus twice of Ni. So this is a net cell reaction. You'll realize that this net cell reaction, you don't have to really multiply, okay. So you simply put 0.4 plus 0.6 and you'll end up getting your answer, which is equal to one electron, one volt, okay. So this is one common mistake that you end up multiplying the electrode potentials also, which is not really required, okay, so that's one. In all the energies, you definitely do. When you do delta H, when you're looking at delta E, internal energies, all of those there, it is fine, but not in electrode potential, that's point number one. Then the second point number two is that, I also find that students are confused in terms of really writing the cell reaction. Now please note, firstly, this is the way. The E0 of cell is nothing but E0 reduction, okay, minus, so this is the reduction potential of the reduction cell minus the reduction potential of the oxidation cell, okay. So always remember, this is the reduction potential minus the oxidation potential and these are the reduction potentials only, okay. When we write the oxidation potentials, it completely becomes ULTA. So it is E0 oxidation potential of the oxidation cell minus E0 oxidation potential of the reduction cell, okay. This is very important. Please note I'm repeating one more time. E cell is nothing but reduction potential of the reduction cell, reduction potential of the reduction cell minus reduction potential of the oxidation cell, okay. Now, so these are a few important points that you have to really remember. There's a quick question here. So the question is, since atmospheric pressure is almost constant, therefore such reactions may involve the changes in volume. Yeah, so this is a thermodynamic question. I think we'll take it back, okay. Yeah, so the thermodynamic question we'll address. Is there any question on electrochemistry? I would like to answer the electrochemistry questions at this point so that, you know, I mean, we'll just stick to the topic and then I'll address the other questions as well. Anyone who has any other question? I have not been able to see yet. Okay, guys, now I think, you know, we are ending up the closure of, we just have one more, maybe one or maximum two more sessions left. Organic chemistry is what we'll be doing in the next class. You know, I would be very happy, you know, if you guys have any questions, you know, from any chapter, I would like to address them there. Organic chemistry is a very important lengthy chapter. You know, we'll be looking at all the aspects of organic chemistry going forward. You know, I would like you guys to really read the chapter and come when we do the revision. Please, you know, these revision sessions are, will be of no use if you are not really knowing the chapter, you know, before coming to this session, you know, because these revision sessions are more of a, you know, review kind of a thing. So I expect that, you know, you guys actually must have at least read one or two, the chapter one or twice, months or twice, which will actually help you in understanding what we are saying here. Okay, right, so, okay, okay. So I don't see that there are any more doubts, you know, I mean, there are some other doubts that people have, but that's from other chapters, which I'll address. Okay, good. So we will pause here then. I don't see, you know, any, yeah. I'm always available in WhatsApp, you know, there are students who are reaching out. If anyone of you has still not, you know, connected and you feel still feel need help, please make sure that you buzz me up and take, you know, whatever is needed, get your doubts cleared. Do not shy off. This is the last time that, you know, I mean, a couple of weeks and you will be facing your boards. I very strongly urge that if you feel that there is anything that we can do, you know, you're most welcome. You know, I've been having sessions with individuals with maybe a group of two students, three students also, making sure that the doubts are getting cleared. So feel free to just buzz up, share your doubts on WhatsApp and I'll be happy to really look at these, okay? So we'll pause here for today then. You know, I'll see you next on, I think Wednesday we are meeting up again the day after tomorrow with the next topic, Wednesday or Thursday, whenever it is next, according to your time table and I'll help you understand organic chemistry then. Feel free to really look at these different aspects and, you know, I'll see you then, okay? Thank you so much. All the best for your studies. Stay connected and, you know, we'll keep on solving problems. Thank you, bye-bye.