 Electrolysis of aqueous NaCl is one of the important processes using which NaOH is produced in industries. But before we go to the case of aqueous NaCl, let me just quickly recap how electrolysis of molten NaCl takes place. We have seen this type of an electrolysis setup before, where we have these two electrodes connected to a power source and the electrolyte here is molten NaCl, which we get when we take some NaCl and we heat it until it reaches this molten state and we get Na plus and Cl negative ions. And in this case, if we write down the reactions that take place at the electrodes, on the cathode side, we have the Na plus ion gaining an electron to become sodium metal and on the anode, we have two chlorine ions coming together to form chlorine gas. And in this case, if we want to write the cell reaction, we multiply this reaction by 2 and add it to the reaction below. And so we get this as our final cell reaction. Now as you can see, when we performed electrolysis of molten NaCl, we got sodium metal and chlorine gas as the products. But what if we take another cell in which the setup is almost the same, except this time, the electrolyte is aqueous NaCl, which is also called brine, and we perform electrolysis here. And in this case, when you look at the products, we see that the products of electrolysis are chlorine, which was same as in the case of molten NaCl, but here we also have two other products. We have hydrogen gas and NaOH. So then the question is, in molten state, when we got sodium and chlorine gas, why is it that by using aqueous solution, we're getting these different set of products? And one clue why this could be happening is this aqueous here. So these products here have got something to do with the aqueous nature of the electrolyte. So let's see what's happening here. Let's consider the cell here, which is the same as before, and with the electrolyte as aqueous NaCl, which is a solution of water and NaCl. Now in this case, let us look at the reactions that take place at the electrodes. Let's start with the cathode. So the cathode is connected to the negative terminal, and so you can think of the sodium ions moving towards it, gaining an electron, and forming sodium metal. This is one possible reaction, and this is also same as what we saw in the case of molten NaCl. But in this case, since we've used aqueous NaCl, and there is water here, there is one more possible reaction that can take place here, and that is that water can gain electrons and get reduced to form hydrogen and give hydroxide ions. So what's happening at the cathode is that both of these reactions are sort of competing. So if we want to find out which of them will actually happen at the cathode, we need to go to our thermodynamics and check which of them is feasible. And for that, what we do is we can look at the reduction potentials for both of these reactions. So in most textbooks, you will have somewhere in the appendices, the standard reduction potentials and the reactions. So I've picked this data from there, and we also know something about the standard reduction potential. The higher the value, the more the chances of reduction. And since these are negative numbers here, the standard reduction potential in the second case is larger than that in the first one. And so we know that at the cathode, the second reaction takes place. So this reaction right here is our reaction that takes place at the cathode. Now in a similar manner, let's look at the reaction at the anode. Again, just like before, the anode is connected to the positive terminal and it will attract the chlorine ions. And one possible reaction is that chlorine ions come together to form chlorine gas. But again, just like before, water also can get oxidized here. So another possible reaction is water getting oxidized to give oxygen and H plus ions. And just like before, if we think of these as competing reactions, to find out which one will actually occur, we go back to our reduction potentials. Again, I've noted down the reduction potentials for these reactions. But the point is that these are oxidation reactions. So in this case, higher the reduction potential, higher the tendency to get reduced or lower the tendency to get oxidized. So from thermodynamics, we know that the value with the lower potential is the one that will take place. So we might expect that this is the reaction that takes place at the anode. But if we actually perform the experiment, we see that at the anode, chlorine gas is formed. So then the question is, why does this happen? We've seen before how electrolysis can be affected by the change in composition of the electrolyte. And we used the Nernst equation to describe how the cell potential changes with temperature and also concentration, which comes in in this term Q, which is the reaction quotient. So now for this reaction, if we apply Nernst equation, let's say we want to calculate the cell potential. So we know the standard cell potential, which is this 1.23 volts. We can take the temperature to be 298 Kelvin or 300 Kelvin. And instead of the Q here, we can write this to be equal to the concentration of H plus ions because the oxygen and the water are in their standard states. But we know one more thing about this solution. So the pH of this solution is 7. And we know a relation, which is pH is equal to minus log concentration of H plus. So using this relationship, we know that the concentration of H plus ions in the solution is equal to 10 to the power minus 7. So what if we substitute this here? In that case, the E cell will be equal to E naught, which we know minus RT divided by 4 times F because 4 is the number of electrons that are involved in this reaction times ln 10 to the power minus 7. Now, if we use the property of logarithms and bring this minus 7 down here and multiply it, we see that the sign of this expression changes. This becomes a plus 7 times this value. So we actually don't need to calculate all of this. But if we do, what we'll find is the effective cell voltage is higher than this 1.23. In fact, depending on the temperature and the pH, which can slightly vary, the actual cell potential can be up to 1 volt higher. So instead of this 1.23, it could be as high as 2.23 volts. So this additional potential is called overvoltage, which comes in because of these external factors like variation in the concentration of the electrolyte or the temperature. In case of chlorine also, there may be such variations. But even then, what we see is that the reduction potential for water will be much higher than that of chlorine. And because of this, as we saw before, the potential value for chlorine becomes the lower value of reduction potential, which means it has a higher tendency to get oxidized. And which is why we see that this reaction is the one that takes place at the anode. So if you look at the reactions for this cell at the cathode, we have water getting reduced to give hydrogen gas and hydroxide ions. And at the anode, we have chlorine gas, which is released. So we can put these together and write the cell reaction. In this case, when we have used aqueous NaCl. So if we write down the reactions occurring at the cathode and anode, in the case of aqueous NaCl, and if we add them up, we get this cell reaction. And you can see that the products form are chlorine, hydrogen and this hydroxide ion. And because there are Na plus ions in the solution, we can write this as NaCl and this as NaOH. And so what we see is that unlike the molten case where we had sodium and chlorine as products, when we use aqueous NaCl, we get chlorine, hydrogen and NaOH.