 Okay, shall we start now? See, ionic equilibrium is like, you know, the last chapter of physical chemistry of grade 11. We have three, four, five more chapters after this. Okay, in this one, in physical chemistry, if you see, that is there in grade 12. This is the last chapter of grade 11 physical chemistry. Yeah, it is a different chapter in NCRT. It is given in the under one chapter only. That is equilibrium. They have given chemical and ionic both together. But this is a different chapter, chemical equilibrium we are done. Okay, now we are discussing about ionic equilibrium. Like I said, it is the last chapter of physical chemistry of grade 11, right? Important one, for example, point of view and a little on the tougher side, you can say. Because, you know, basic questions you will get, you will do that. But there are a scope that you can get very tough questions from this chapter also, okay? If you look at the vastness of this chapter, it is like the two most toughest chapter if you consider in grade 11 physical chemistry, can take off thermodynamics and ionic equilibrium, okay? So relatively tough, you can say, from the other one. Like it's not like ionic equilibrium, which is quite easy and easily you can understand all these things, right? So here we have again conditions, you have to, you know, analyze the condition and then accordingly you can use the formula. That kind of analysis is required over here. If you talk about ionic equilibrium, so here we'll discuss about the equilibrium among the ions in the solution, right? Equilibrium among the ions in the solution, what kinds of equilibrium is there? So, and what is the property of that? How the equilibrium exists, okay? So this chapter, mainly we are going to study about, study about acid base. We're going to study about acid base in this chapter. We are going to study about, you know, pH scale, how to calculate pH that also we'll see. And after this, we have again, neutralization reaction and all hydrolysis, solid hydrolysis buffer solution. There are so many things. Mainly the entire chapter will go around acid base and its pH scale, okay? Acid base and its pH scale. That is what the entire chapter is, okay? So what I said, ionic equilibrium deals with the equilibrium among the ions, right? Equilibrium among ions, correct? So when you dissolve a compound in any solvent, right? It may form ions over there and then the equilibrium of ions we are going to discuss over here. Like for example, if you take NaOH, solid dissolved in water, H2O, it converts into Na plus aqueous and Cl minus aqueous. So what happens, you have a container. In this container, you take water, we take water and we add NaOH over here. NaOH you are considering only one particle, but in this solution here, it is present at two particles, like two ions are there. Na plus OH minus, correct? So it furnishes ions into the solution. It's not like all the compounds will furnish ions into the solution. Like suppose if you have this one, you have water and you add, suppose you add sugar into this. It is here sugar only. It won't furnish ion over here, right? It gets sugar only. So we have basically two types of compounds. You have basically two types of compounds over here, right? One type which provides, I have taken NaOH only here, no? So here also I've taken, see what I have written. One second, I'll just correct it. NaOH is fine, and here we'll have, here we'll have OH minus. Okay. Sugar if you add, sugar won't furnish any ions into the solution. It is here dissolved in the solvent like this, okay? So we have basically two types of compounds. The one which can furnish ions into the solution and the one which does not furnish any ion in the solution. The compounds which furnish ion in the solution or any solvent, we call it as electrolyte. What? Electrolyte. This kind of compounds only are responsible for the electrical conductivity of any solution. Ions can carry charge over here, right? You know in metals we have electrons at the, you know, charge carrier. Here we have ions which carries charge and hence it shows electrical conductivity, right? This kind of solution. This kind of solution is non-electrolyte. It does not show conductivity. This kind of compounds is non-electrolyte. Non-electrolyte are those compounds which does not furnish ion, which does not furnish ion in the solution. Did you get this? Yeah, it's one. Since we talk about the equilibrium of ions, so we must have ions present in the solution. So here in this chapter, we mainly deals with electrolytes. Our focus will be on electrolyte only because it is the one which can produce ions in the solution. So we'll be considering mostly electrolytes over here. Clear? So right down the next heading here, we'll talk about little bit of electrolytes and they'll be move on. Heading all of you right down electrolytes. These are the compounds, compounds which can produce ions in aqueous medium, in aqueous medium are called aqueous medium simply. Compounds which can produce ions in aqueous medium, they have capability of, capability of or capability to produce electricity. Example of electrolytes, we can have many example, any acid we can take, any base we can take, or any salt also we can take. All these comes under the category of electrolyte. Clear? Now you'll see electrolytes are classified into two categories. We have basically two types of electrolyte. The first one is strong electrolyte, cover this down, right down there are substance, there are substance which completely dissociates in water, completely dissociates in water or any polar medium, or in any polar medium, complete dissociation is there. If you talk about degree of dissociation for this kind of substance, degree of dissociation alpha equals to one because we have complete dissociation. So the degree of dissociation for a strong electrolyte is always one. Okay, got it? Done, all of you copied guys? Okay, next we have the second one, which is a weak electrolyte. We'll be dealing with weak electrolyte mostly in this chapter. Write down the other substance which does not dissociate completely, substance which does not dissociate completely and goes under partial dissociation, goes under partial dissociation. Example of strong, I did not give you, I'll write down it. Example of weak electrolyte, you write down. You can have weak acid. Any example of weak acid could you give me? We can take weak acid, we can take weak base, some salt also we can take over here. Some salt, that is AGCL. These are the weak electrolyte. CS3COH you can say Shradha, not NA. NA is a salt, CS3CON is a salt. CS3COH, yeah. Organic acids are generally weak acids, okay? Inorganic are strong except H2CO3. H2CO3 is a weak inorganic acids. If you take the example of strong electrolyte here, it will be a strong acid. I'll write down in short here. A strong acid, we can have strong base or we can have some salt often. Generally the salt of strong acid and strong base are called strong electrolytes. So for weak electrolyte, the degree of dissociation, alpha, alpha is less than one. Start it, equal to one or greater than one. Alpha is less than one over here. Okay, now you see this. Suppose we have a weak electrolyte HA and it dissociates like H plus plus A minus. So the initial concentration of this acid, HA is C, this is zero, this is zero, initially. At time T is equals to T, it's obviously some part of this, HA will convert into X plus and A minus. So it is C minus, C alpha, C alpha and C alpha. Where alpha is the degree of dissociation, C is the concentration we have. Did you understand this? Yeah, tell you guys. Any doubt, no? So what is KC here? KC equals to concentration of H plus, concentration of A minus, concentration of HA, okay? H plus and A minus is C alpha square, C minus C alpha, which is C alpha square divided by one minus alpha. Alpha is the degree of dissociation. So, copy this down first. See, HA is a weak electrolyte, so its degree of dissociation is very small, very less. Very small amount of HA will convert into H plus and A minus. So if we take one minus alpha, if alpha is very small, very small, then one minus alpha is almost equals to one, right? And then this becomes, we are neglecting this alpha with respect to one actually, approximation. And then this becomes K is equals to, K is equals to C alpha square and alpha is equals to K by C root over of it, right? So now you see here, what we can say, concentration and degree of dissociation is inversely proportional, right? So we can conclude from this on dilution, on dilution of weak electrolyte, on dilution of weak electrolyte, the degree of dissociation, the degree of dissociation increases on dilution of weak electrolyte, the degree of dissociation increases. What do you mean by dilution? So in this, if I write down in this relation here, as concentration decreases, alpha increases, right? And it was also it's true, as concentration increases, concentration increases, then alpha decreases because both are inversely proportional. This relation of concentration and alpha, this relation we call it as, we call it as host-wards dilution law, host-wards dilution law valid only for weak electrolyte, valid only for weak electrolytes, because for strong electrolytes, for strong electrolytes, alpha value is already one, right? So there's no point of calculating all this. So it is host-wards dilution law, it says on dilution, the degree of dissociation that is alpha increases, copy this down, okay. So you see, alpha is keep on increasing, right? As C decreases, because dilution means what you're diluting the acid. You keep on adding water, that means strong electrolyte, Kc is equals to infinity, how? We'll take alpha is one, how is Kc is infinity for strong electrolyte? So here what happens, you see, as you keep on diluting the acid, dilution means addition of water, as you keep on diluting the acid, right? It's alpha is increasing, means it's degree of dissociation is increasing. Alpha suppose, alpha suppose it is less than one initially as you keep on diluting, alpha is increasing slowly and slowly, right? And it is going closer to one, right? Tends to one, we have no. And once it tends to one, it starts behaving as a strong acid, as it has a tendency practically, it's not possible, but it has this tendency that as you keep on diluting acid, its degree of dissociation increases and goes to a maxima, right? Where it starts behaving as a strong acid. So right down one last point over here, on very low concentration, on very low concentration, weak acid behaves as a strong acid. On dilution, weak acid starts behaving as a strong acid because alpha is keep on increasing, it tends to one. Done all of you? Keep always this thing in mind that post-wall dilution law is valid only for weak electrolyte, right? And in this chapter, mostly we'll be dealing with weak electrolyte only the alpha value, correct? Okay, now if you talk about acid and base, see for acid and base, we have three different definitions for acid, okay? Three different different form, this thing define our definitions. So we'll see all these definitions over here. Like I said, we have three different definitions of acid. Acid can be defined as, once again, we can have Arrhenius acid, then we have Bronsted acid, and the last one is Lewis acid. What is Arrhenius acid? See, Arrhenius acid are those substance which releases H plus ion in water. Just a second, okay. So Arrhenius acid you see, these are the substance which increases H plus concentration in water. H plus means what protons? Bronsted is the acid which increases the concentration of H plus ion, of H plus ion. Here we don't have the constraint of solvent. Here the solvent is water. Here we can have any solvent. Lewis acid is the electron pair acceptor, electron pair acceptor, or Lewis acid. The one which takes electron pair are called Lewis acid electron pair acceptor, right? Arrhenius and Bronsted have the similar kind of definitions. In Arrhenius what we say, it is the proton donor in water, the solvent must be watered. Bronsted we can have any solvent. We don't have any constraint for solvent to it. Copy this, okay. One more thing here you must understand that in this Bronsted acid definition, we'll talk about this proton. That is we have proton donor it is. Here we have H plus donor in water. So the definition is what? The substance in donate H plus only in water is Arrhenius acid. The one which donates proton in any solvent is called Bronsted acid, okay? We don't have the constraint of solvent in case of Bronsted acid. Now if you see here, we also know acids are those substance which turns blue lit must to red, okay? That also you write down. These are substance which turns blue lit must to red, okay? Now, suppose we have a compound. If a compound contains hydrogen, it does not mean that all the hydrogen are acidic hydrogen, write down. If a compound contains hydrogen, hydrogen, then all hydrogen may or may not be, may or may not be acidic. Like for example, you see, suppose we have CS3 COOH. So we have four hydrogen present over here, but this one, the one which is attached to the electronegative element is the acidic hydrogen, yeah? All hydrogens are not acidic. It is the one which is attached to the electronegative atom. If you talk about H3PO4, H3PO4 the structure is this. OH, OH and OH. So all three hydrogens are acidic hydrogen here. Important, clear, done. One more understanding you must have. If you look at the definition of Brownstead and Arrhenius acid, okay? Arrhenius acid, the solvent must be watered. Brownstead acid, there's no condition of solvent we have. So we can say Arrhenius acid, all Arrhenius acids are Brownstead acids, but all Brownstead are not Arrhenius acid. Can we say that? Because Arrhenius acid, the solvent is watered. Brownstead concept does not have any constraint with solvent, right? It says whatever the solvent it is, if it is donating hydrogen atom, then it is a solvent, right? This also you must take care of the difference between the Brownstead and Arrhenius acid. Similarly, we have three definitions of base also. Arrhenius base and then we have a Brownstead base and then we have a Lewis base. Arrhenius acid is H plus donor. Yes, it's the same thing. Laurie Brownstead acid, Laurie Brownstead base. Same thing it is. Laurie Brownstead acid is same as Brownstead acid. What is the definition of Arrhenius base? Arrhenius base are the compounds which can donate OH minus ion to water, donates hydroxy ion in water. Brownstead base is the one which accepts H plus, Brownstead base are the compounds which accepts H plus, okay? Lewis base are electron pair donor, electron pair donor. This is three definitions we have of Arrhenius base, Brownstead base and Lewis base. Clear, no doubt, tell me. Do you know what is pH? What is a pH scale? So pH scale is we use to measure, to measure the acidic behavior of a compound, right? And it is defined mathematically as minus log of H plus concentration. Since if you substitute here, minus log of H plus, if you find out this minus log of H plus, you'll get the pH of the solution. What is the pH of a neutral solution? Seven, why it is seven? Can we have any other value for neutral solution? Possible? As a first of all, you tell me, why it is seven? Any clue, any idea? Midpoint of what? Yeah, that's what my question is. Why it is one to 14 scale? Why it is 14? Why not 16? Why not 13? Why not any other value? Okay, that is what we are going to understand over here. Second thing, we know the acidic solution, pH is less than seven, correct? So why pH is less than seven? And for base, why it is more than seven? That is also we are going to understand, okay? Under any condition, is it possible to have a neutral solution of pH six or 6.5 or 55.5? That is also we are going to understand, okay? So let's see this. See, if you talk about the dissociation of water, dissociation of water. So we have H2O, H plus and OH minus. Like this, it dissociates, okay. Okay, we'll do this and then we'll take a break. Give me five minutes, okay? Yeah, give me five minutes. H2O dissociates like this. Now, if I write down the equilibrium constant here, K, this is equals to H plus into OH minus divided by concentration of H2O. Concentration of H2O does not change much, okay? Why? Because this H2O is a polar covalent compound. Covalent compound, the dissociation is very less, is small or very less, right? Dissociation is small. Hence we assume the concentration of H2O is almost constant. It's not changing much, so it is a constant. So what we'll write next, you see, we'll write the equilibrium constant K into concentration of H2O is equals to H plus and OH minus concentration. This product, we call it as KW is equals to H plus and OH minus concentration, correct? This KW is called, KW is called the ionic product of water. This we call it as KW, write down all of you. It is dissociation constant, dissociation constant, ionizational constant, ionizational constant. And we also call it as ionic product of water, of water. Now, this value of KW, we have determined experimentally, okay? And it is temperature dependent because KW, you see, it is a function of equilibrium constant. So it is also temperature dependent. It depends upon the temperature. So at 25 degrees Celsius, the value of KW is find out to be 10 to the power minus 14, this you have to memorize. If you find out it at 80 degrees Celsius, KW equals to 10 to the power minus 14. This is the value of KW we have at different, different temperatures. You have to memorize this one. This one is important. One more thing you see, the dissociation of water here is an endothermic process. Let me write down here. Dissociation of water is endothermic process. An endothermic process, tell me, if temperature increases forward or backward, forward or backward, temperature increases forward or backward. Endothermic temperature increases. I think it is forward, is it? Yeah, forward, Gayatri? Is it forward? Yes, it is forward. So you see as the temperature increases, this goes in forward direction. So concentration of H plus and OH minus also increases. And when the concentration of H plus and OH minus increases, the value of KW should also increase. If this and this will increase, then value of KW should also increase. That's why you see at 25 degrees Celsius, the value of KW is lesser than the value at 80 degrees Celsius. Yeah, understood. So write down those two lines. KW is temperature dependent. Temperature increases. KW increases as temperature decreases. KW decreases. We haven't done this. Okay, we'll continue with this after the break. One second, guys. Let's don't go away. One small thing I wanted to ask for those like for the bio students, are you comfortable for this week if we have a session?