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We've seen a reasonable number of acid reactions and base reactions. So let's just write down a few of them for review and let's see if we can see a general pattern here. And a lot of this might not be any news to you. So if we have hydrogen flouride, or if it's in an aqueous solution it's hydrofluoric acid. We know that this is a weak acid-- it doesn't disassociate completely. So it's in equilibrium. Some type of equilibrium. That doesn't mean the concentrations are equal. This hydrogen disassociates. Actually, we know in reality, it tags a ride along with another water molecule and forms hydronium. And then you have-- and, of course, this is still aqueous. Everything is going on inside water. And then you have left over your fluoride anion, or negative ion. And that's also in an aqueous solution. And we could have rewritten-- actually, let me write another reaction here, just so you can see the general pattern. Let me write another acidic reaction. Let me write ammonium. So that's NH4 plus-- it's ammonia with an extra hydrogen in an-- let me just write that-- in an aqueous solution. That can disassociate to one of those hydrogens popping off in an aqueous solution. And then you have ammonia, NH3. That's also in an aqueous solution. Now both of these describe an equilibrium reaction, but it kind of implies that we're dealing with weak acids. You take an acid, and they're producing hydrogen, which is at least the Arrhenius definition of an acid if you looked at the Bronsted-Lowry definition where they're donating protons to the solution. They're creating hydronium, they're donating protons to the water around it. But it kind of describes it as an acid. But we know it's a weak acid, so this reaction goes in two directions. So we can write the same reaction essentially as a basic reaction. So we can-- instead of saying hydrofluoric acid is our acid-- we could say hey, if I just have a fluoride anion, if I just have a negative fluoride here, I could say a negative fluoride, if I put that in-- actually, I keep making that same mistake. The fluorine does not have an l in it. I do that because chloride does. Let me erase this. Hydrogen fluoride is HF. So it's just F there. Let me go to the periodic table. See, I always confused fluorine with chlorine because F is just for flourine. But you get the point. Ok. So I could rewrite the same weak acid equilibrium as a weak base equilibrium. Or I could say a negative fluorine anion in an aqueous solution is in equilibrium with-- and now we're saying I'm considering this a base, which means that it's going to increase the concentration of OH. So what this might want to do is it might want to grab some hydrogen from some of the water that's in the aqueous solution. So it grabs some hydrogen and becomes hydrogen flouride, or hydrofluoric acid. Let me do that in that magenta color. It's aqueous. And where did it get this hydrogen from? Well, it got it from one of the surrounding water molecules, which was H2O. Since it gave away one of the hydrogens now it's just OH minus. So the surrounding water molecule is OH minus aqueous. Now these might look different. This is donating a hydrogen to the surrounding medium, and then you're left with just the fluorine molecule. Well, this is essentially creating a hydroxide molecule out of the surrounding medium so it looks basic, but if you think about it, these reactions are the same. I mean, you could have just gone in reverse direction. You could say, hey, this is going to react with some random hydronium molecule or some random, free-standing proton out there. And then it could form hydrogen fluoride, but we know that hydronium isn't just sitting everywhere, that whenever you take the reverse reaction, whenever you're going in this direction this doesn't have to grab this hydrogen from an H3O. It could grab it from an H2O. And then you would have this reaction. These are equivalent. And we could do the same thing here for ammonium and amonia. We could write ammonia as a base. NH3 is in equilibrium as a weak base with-- it can grab a hydrogen from its surrounding medium and become NH4 plus in an aqueous solution. And then it would have grabbed that hydrogen, probably from a water molecule because that's what's around it. And so that water molecule will become an OH minus. And so now this looks ammonia is a weak base. Ammonium is a weak acid. But these are equivalent reactions. Now, you're probably already seeing a relationship here. Ammonium is a weak acid. Ammonia is a weak base. And what's the difference between the two? Just an H. Hydrofluoric acid is a weak acid, just a fluorine anion. A negative fluorine is a weak base. And what's the difference between the two? They're just difference of a hydrogen. Let me write that down. So let me write weak acid. And then you have your weak base. So your weak base-- let me write our weak acid is first. We had hydrofluoric acid, and then the weak base is when you essentially just dump the hydrogen, just the hydrogen proton. You kept the electron, so that's why it's a negative charge right there. Hydrogen without its one electron is just a proton because it has no neutrons. The other one was NH4 plus. You dump one of the hydrogens and you get NH3. So what's the difference going on? These are all minus a hydrogen. Or if you go this way, you're plus a hydrogen. So you have these kind of conjugates. And this has all been a long-winded way of introducing you to this idea that you have these conjugate pairs. Like hydrofluoric acid, or hydrogen fluoride and just the fluorine anions. So these are conjugate pairs. Which are essentially two molecules that are identical except for a difference in one hydrogen. No more than one hydrogen. One day there might be a test where someone shows you two molecules that are separated by two hydrogens-- those would not be conjugate pairs. For example, if I show you H2O and OH minus, these are conjugate pairs. Because this over here is exactly this minus a proton. And, in fact, let me be clear that it's not just minus the hydrogen-- minus the proton. One of them is keeping the electron. So this is minus a hydrogen proton, this is plus a hydrogen proton. So the difference between these two are just a hydrogen proton. So these are conjugate pairs. Now, if I were to say that H3O and OH minus, you might be tempted to say, hey, this is very acidic, this is a base, this is a conjugate pair. But no, there's a 2 H, 2 proton difference. This is H3O plus. There's a 2 proton difference, so these are not conjugate pairs. So let me just cross that out. But these are. Now, you could say, if you have H3O, you might say, hey, what's the conjugate base-- and that's a new word I just introduced you to-- what's the conjugate base for H3O if H3O is an acid? Well, you take one H from it and you get H2O. So this is a conjugate pair. And I just said a word without defining it, so now let me define it. Within every conjugate pair you have an acid and a base. And if you say, oh, what is the conjugate base for hydrofluoric acid, you get rid of a hydrogen and you say, oh, it's just this fluorine anion. If you said, I have some of ammonia, as a base, what is its conjugate acid? So if someone asks you, what is a conjugate acid, you add a hydrogen proton to it and you get ammonium. So I could call these the conjugate acid. Let me just change terminology-- conjugate. And we'll see that actually, you don't have to be using a weak acid or a weak base. Conjugate acid, and then you have a conjugate base. And even though something might be a conjugate acid or conjugate base, it doesn't necessarily mean that they're very basic, for example, or very acidic. If I have hydrogen chloride, we know this is a strong acid. Hydrogen chloride. Its conjugate base, we essentially just get rid of one of these hydrogen protons-- it doesn't take its electron with it. So it's just going to be the chlorine negative ion. This is its conjugate. If I gave you a chlorine negative ion and said, what's its conjugate base-- what's its conjugate acid? You'd say it's hydrochloric acid. If I gave you hydrochloric acid and I say, what's its conjugate base, you get rid of a hydrogen proton only, and you're left with the chlorine negative ion-- You said that its conjugate base. Now, with that said, we know that when you put hydrochloric acid, we know this reaction. This was I think the first reaction we looked at in aqueous solution. It disassociates completely to form hydrogen protons plus chlorine anions-- everything, of course, in an aqueous solution. Now, the fact that it disassociates completely, that this is not an equilibrium reaction, this implies that this guy is more basic than water. He has no temptation-- no, no-- let me say that he is less basic than water. He has no temptation to grab these hydrogen protons from the surrounding medium to reform hydrochloric acid. This reaction does not go in this direction. So even though this chlorine anion, or negative ion of chlorine, is a quote unquote conjugate base of HCL, that doesn't necessarily mean it's that basic. This is less basic than water. It wants the hydrogen protons less than, let's say, hydronium. So if you put some chlorine plus some hydronium, or let's say you have some H plus, you're not going to reform hydrochloric acid. So this is not really basic even though it's considered a conjugate base. And that's generally the case whenever you're dealing with strong acids, like in the case of hydrochloric acid. If I had a big solution of just chlorine anions in water, so I just had tons of a super high concentration of chlorine anions and water, because it's not going to do anything to change the actual hydrogen or hydroxide concentration in the water because it's less basic than the water itself-- it doesn't want to take or give anything to the water-- the PH, so if you have a soup, the PH would be 7. If you have chlorine minus in an aqueous solution-- and I don't care what its concentration is, you could have 10 molar of it-- the PH is still going to be 7 because it's not going to change its solution. It's not going to change the PH, just this by itself. Obviously, if you put hydrochloric acid in an aqueous solution this will change it, because you're going to be dumping all of these hydrogen protons into the solution. So in general-- I mean, you can kind of remember it, but I think it's maybe common sense-- a strong acid's conjugate base is neutral in water. Neutral. So that means no impact on PH. You say chlorine plus H2O, I mean, you're essentially still going to have chlorine plus H2O. You're not really changing the concentration. Now, on the other hand, when you're dealing with weak acids, so this reaction will go in the other direction. If you put some fluorine in water, it will grab some hydrogen-- not necessarily a ton of it, but it will grab some hydrogen from the surrounding water-- and increase the hydroxide concentration. It's increasing the concentration of this thing right here. So it is making, in this case, it is making the solution more basic. It's increasing the PH of the solution. So whenever you have a weak acid its conjugate base will be a weak base. And you could make this statement the other way around. If you have a conjugate-- let me switch colors, this is getting annoying-- conjugate base-- sorry, a weak base, its conjugate acid is going to be a weak acid. So hopefully you get the idea here. It's actually not that fancy of an idea. It's just that if you have an acid, its conjugate base is just that acid minus a hydrogen. If you have a base, its conjugate acid is just that thing plus a hydrogen. Actually, let me just do a bunch of problems here just to really hit the point home of what we're talking about. Let's just do a bunch of them. So if this is the acid and this is its conjugate base, so if I have-- I mean, you don't even have to know the words. If I have that, the conjugate base, well, I'm just going to get rid of a hydrogen proton. So NO3 minus. I didn't get rid of the whole neutral hydrogen molecule. Remember, I just took a proton away, the electrons stay the same, so I have a negative charge. Let's say we did H2SO4. So its conjugate base, get rid of a hydrogen. HSO4 minus. If I have hydrogen bromide, get rid of a hydrogen. It's BR minus. And this is a strong acid. So this is going to be a neutral-- if you put this in water, it's really not going to do anything even though you are calling it hydrogen bromide's conjugate base. Now, if we go the other way. If we give you the base, if I give you OH minus, what's its conjugate acid? Well, you add a proton to it, you get H2O. If I have H2O-- we already did that-- you add a proton to it, you get H3O plus. If you have-- I mean, we could just keep going. Let's say I have that. If I add a hydrogen to it, I have H2. There you go. And it's neutral now because I added a proton. Anyway, hopefully I haven't beaten this horse to death and you understand what conjugate acid and bases are all about.