Understanding Acids and Bases in Chemistry

All right NINJA Nerds, so what we're gonna do in this video is we're gonna talk about acids and bases. So we're gonna go through, first off before we get into any calculations, because that's what we're gonna spend a lot of time on. Kind of the key to chemistry and getting better at it is doing tons and tons of practice problems. And so that's our goal. What we're gonna do here at NINJA Nerd Science is help you as much as we can with a lot of practice problems. Before we dive in, let's go ahead and get some just baseline introduction, introductory material out of the way. So first off, let's go ahead and define what an acid and what a base is. Alright, so we're going to do that, then we're going to talk about all the strong acids and strong base. We're going to talk about conjugate acid, conjugate base, pH scale, and we're going to talk about this KWKAKB, PKPKB. Alright, so let's go ahead and dive in first with definition of an acid and a base. In order for us to do this though, we have to go through three different types of definitions. Alright, so the first one we're going to talk about is Bronsted-Lowry. Okay, so Bronsted-Lowry. So, Bronsted-Lowry acid. When we're talking about a Bronsted-Lowry acid, he said that when you take an acid and you put it into a solution, an acid gives up protons. So, how does he define an acid? He defines an acid as a proton donator. Or we can even draw this like this. He loves to give up H plus ions because we're gonna see this a lot, okay, whenever we're doing a lot of these practice problems. So again, what is a Bronsted-Lowry acid by definition? A proton donator. Well, if an acid is a proton donator, someone has to be able to gain that or accept that proton. Who accepts the proton? The Bronsted-Lowry base. So the Bronsted-Lowry base is going to be the guy who is accepting the proton. So he's a proton acceptor. Alright, so that's pretty easy right? So it's not too bad there. So for example if I were to ask you, if I were to draw right here, let's say I did this, H A plus water. and it yields two products, let's say one of them is A negative and the other one is H3O+. If we were to go based on this definition of Bronsted-Lowry, what happened to this HA? It turned into A negative. So what did it do? It gave up a proton. So he must be, by definition, the Bronsted-Lowry acid. And then what happened to this H2O? He turned into H3O+, hydronium. So he must have accepted that proton. And if he accepted the proton that the HA gave, he must be the base. Now while we're here, let's just go ahead and get the... terms out of the way since we're already here conjugate acid conjugate base let's go ahead and get those terms out of the way because we're already here so an acid when he gives up the proton so he donated that proton to the H2O he turns into a different term now that we can call So he was an acid, but now he's a base. By definition, if he's a base, usually bases have negative charges, right? So they're really good at having negative charges. That's kind of the way that we can see bases is they're negatively charged molecules or atoms. So we're going to say that this is a base, but it's another type of base. It's the conjugate base. So this acid, it gets converted into a negative, which we're going to call the conjugate. Base. And again, why is he the conjugate base? Because the acid gives up a proton and turns into a conjugate base. Well, if he turns into the conjugate base, then this must be the acid. Why is he the acid? Because this water, when he accepts that proton because he's the base, he turns into an acid. So he is the conjugate acid. So if we were to combine the pairs... This acid and this conjugate base, this is the acid conjugate base pair right there. And this right here is the base conjugate acid pair here. Okay? So not too bad there, right? So that's just giving you an understanding of Bronsted-Lowry, what an acid and what a base is. And then what is an acid, what is its conjugate base, what is a base, and what is its conjugate acid. Alright, let's go into the next individual who came up with this, another theory. This guy is... Lewis. So Lewis came up with a theory of acids and bases, but he looked at it in a different way. Bronsted looked at protons, Lewis looked at electrons. So Lewis said if an acid is donating a proton it has to be accepting electrons. That's his definition. He says by definition a Lewis acid is a molecule who does what? Eps electrons, so he is an electron acceptor. Okay? So while this guy is donating protons, he's also accepting electrons. Okay cool. So then a base, what was a base doing? A base was accepting a proton, so he must be donating electrons. So he is an electron donator. Alright. Not too bad here then, right? So this is, if we were to go based off this, we could apply the same concept here with the Lewis acid and Lewis base. There is two exceptions we're not going to talk about. So Arrhenius, he was the other guy that came up with this theory here, right? And he said... That an acid is a little bit more specific than just a proton donator. So Arrhenius said an acid is actually going to be a guy who gives up, this guy has to give up a proton, so he is a proton. proton donator, okay, so he does give up these actual protons, but more specifically, he gives up these protons in an aqueous solution, so we'll say in water. in an aqueous solution. So he even kind of expanded more on that. He said an Arrhenius acid is a proton donator in an aqueous solution. Okay, the definition of an Arrhenius base, right, so an acid is the one who's going to be able to give up a proton produce an H plus, you know, he's producing a proton in an aqueous solution. A base is just producing hydroxides. So what is it doing? It's producing, or we could say even donating, hydroxides. donator if you'd want to say that or it's producing hydroxides in an aqueous solution. Okay so that's the definition of an Arrhenius acid and base. An acid is going to be one who gives up protons in an aqueous solution whereas a base is going to be producing hydroxides in an aqueous solution or donating those hydroxides in an aqueous solution. Okay so now we have an understanding of what Bronsted-Lowry acid, Lewis acid and base, Arrhenius acid and base and we even know what conjugate acid and conjugate base pairs are So let's go ahead and check those off. So what do we do? We did the definition of an acid and a base, and we did what conjugate acids and conjugate bases are. All right. Let's talk about what strong acids and strong bases are. All right. Well, first off, how would you define a strong acid? A strong acid is defined as something that can deprotonate or disassociate completely. So, for example, I'm going to write down all of our strong acids here. So let's go ahead and put them right here. here strong acids these guys are going to be hydrochloric acid, hydroiodic acid, hydrobromic acid, HNO3 which is nitric acid, HCl4 which is perchloric acid. H2SO4 which is sulfuric acid. So how many do we have? 1, 2, 3, 4, 5, 6. These are our six strong acids. And again, what does it mean that they're a strong acid? That means if you guys, if you put this into an aqueous solution, they're going to disassociate so it's a one arrow movement. If you notice I did not have an equilibrium arrow. Where did I have an equilibrium arrow? Here. Here I will not because all of the HCl will deprotonate into the form of H plus and Cl negative. And again, a strong acid is defined as a proton donator. So it's going to give up that proton and it completely disassociates in an aqueous solution. So what would this one look like? HI would give way to H plus and I negative. HBr would give way to H plus and Br negative. HNO3 would give way to H plus. and nitrate. HCl4 would give way to H plus and ClO4 negative. H2SO4 is different. So if you look here, H2SO4, it has an H2. This is defined as a diprotic. acid, meaning he can give up two protons. So if we were to run this reaction, he would deprotonate and we'd turn into HSO4 negative. And then he would react even further and get converted into SO4 2 negative, right? So what we give up, we should actually show the H plus in there. So again, it would react here and give you H plus and this guy. Then this would react even further and produce sulfate and H plus. So again, these are six strong acids and the reason why they're strong acids is because they disassociate completely. Alright, let's do strong bases. bases same thing the reason they're called strong bases is because they disassociate completely when put in an aqueous solution so what are this what are these strong bases these ones are easy because if you take all the group one and group two metals there is a couple exceptions but if you take almost all of those so for example we take sodium hydroxide we take lithium hydroxide we take potassium hydroxide we take rubidium hydroxide We can even keep going, we can even do cesium, okay, cesium hydroxide. We can't do francium though. Alright, that's a big mofo and not a good one to do, right? Alright, now what's another one? We can do group 2 metals. So what are the group 2 metals? The group 2 metals are the ones that are in the middle. metals we could do magnesium isn't red necessarily soluble calcium is too but it's kind of slightly soluble but we'll put that one in there COH 2 we could also do strontium hydroxide and we can could even do barium hydroxide but not radium. Okay so in general these are our strong bases. Magnesium isn't really considered to be a strong base because it's not really soluble. Okay so if we were to kind of look at all of these guys here this would pretty much give us all of our strong bases and what are they strong why the strong bases because they can give up or donate all of these hydroxides all of them will disassociate into solution. Okay so if you were to put all of these guys into a solution I'm only gonna do one here if I were to take sodium hydroxide here, put him into an aqueous solution, what would happen? It would give up or completely disassociate into hydroxide ions. Okay, so that's the definition of a strong base and that's the definition of a strong acid. They completely disassociate. It's not a reversible reaction. One more relationship I want to develop here because this is really important when we're going through a lot of these calculations. There is a relationship here and it says the stronger The acid, whenever the reaction occurs, the weaker its conjugate base. The weaker its conjugate base. And this would be the vice versa if you start off with a strong base, the weaker its conjugate acid. It's either way. Let me give you an example here. If I take something like HCl, I react this in water, what would I get? I would get H+, which is a strong acid. acid and I'll get chloride. Well this by definition is a strong acid. This right here is a very very weak conjugate base. That's going to be important whenever we talk about the acidity of and basicity of salts. Okay, alright so again remember that relationship stronger the acid the weaker its conjugate base and like I said you could do the same thing with the other one. You could say stronger the base The weaker its conjugate acid. Okay, so now we have that relationship taken care of. So what have we done so far? We've talked about definitions of acids and bases with Bronsted, Lewis, and Arrhenius. We talked about all the strong acids and strong bases, and we tied in the relationship with conjugate acid, conjugate base. and the relationship of strength. Alright let's get over here to the pH scale. Alright so let's go ahead and talk about pH scale. So pH scale is a is basically a logarithmic scale here and we're going to talk about that a little bit but just know that that it's a scale that goes from 0 all the way to 14 and what it does it helps to measure from 0 all the way to 7 measures how acidic the solution is and then going from 7 all the way up to 14 measures how basic a solution is now again what do I mean by acidic and what do I mean by basic So when we talk about acidic, we're talking about specifically, we're talking about how many protons we have or how many hydronium ions we have. So when we talk about acidic environments, we're talking about the concentration of... of H plus and we put these brackets in there to signify concentration. So the more you go closer to zero, the more acidic your solution is. So if you have a solution that's in between one and zero, that's pretty freaking acidic. So when we talk about acidity, we're talking about H plus concentration. When we're talking about basicity, we're talking about Hydroxide ion concentration. So again, if I were to say, okay, I got a solution that can have H pluses and I got a solution that can have hydroxides. How do I know specifically that it's acidic? Well, that means that the H plus concentration in the acidic environment is greater than the concentration of the hydroxide ion. Alright, specifically in the acidic environment because the more H plus is, the more acidic. In the basic solution, the basic solution we want it to be measured by hydroxide ion concentration. So the greater the hydroxide ion concentration is, then the H plus concentration, that means that we're in a basic solution. What if it's neutral though? Then these two are equal. So then the H plus concentration is equal to the Hydroxide ion concentration and this is an important relationship when we talk about Kw here in a second. Okay before we do that, okay I want to talk about another thing. So I said before that this is a logarithmic scale right. So another words if I wanted to go from zero all the way to five most people would say okay how many people was the pH change well they'd be like oh it was five that's wrong and the reason why is each number you go up from zero to one 1 to 2, 3 to 4, and then 4 to 5 is a 10 times change. It's a 10-fold change. So, for example, if I'm going up here, I'm actually going up what? 10 times Okay, so what is that? 10 times 10 times 10 times 10 times 10. Okay, well that's going to be what? That's going to be a 100,000 fold change. That is a 100,000 fold change. So how much of a change was it? It was a 100,000 times change. So remember that whenever we're saying okay, what's the pH change going from 0 to 5? Well, each change is a tenfold increase So if I go 10 and 10 and 10 and 10 and 10 that is a five-fold increase But specifically by 10 times 10 times 10 times 10, right? That's just a hundred thousand times change there same thing if I was going from 14 to 9 If I'm going all the way from 14 to 9, what am I doing? It's changing by a 10-fold scale. So 10 times 10 times 10 times 10, and then over here to this guy times 10. So same situation, how many was that? We went from 14 all the way to 9, most people would be like, oh, that's a 5 times difference. No, it's a 100,000 times difference, because it's a 10 times each one, okay? Now another thing, because of this base 10, if you will, change, they came up with a way to be able to calculate a specific number which we call this pH, right? And so what they said is, okay, well, we can technically figure out what the pH of a solution is since we have this scale now. And what they said is we could actually say that pH is equal to... We put a negative in there for specifically mathematical sake, and we said it's a log base 10 relationship. It's a logarithmic scale. So pH is equal to the negative log of the H plus concentration. Okay, and that's how we derive this formula to solve for pH. We said it's based on this logarithmic scale, which is a base 10, right? Tenfold change. Same thing, that's pH. So that would give you the H plus concentration. P-O-H. would give you the hydroxide ion concentration. So it's the negative log of the hydroxide ion concentration. Okay, well that's cool because now I have this ability to calculate pOH, I have the ability to calculate pH. What if I add both of these two together? So then I take and I add pOH plus pH. That gives me a specific formula. pH plus pOH, wait, isn't that the whole scale? It is. So it equals 14. And that's how we got a lot of these formulas, right? So what do we have so far? We have pOH equals negative log of the hydroxide ion, pH equals negative log of the hydronium ion, and then pH plus pOH equals 14. In the same way, if you know your algebra, if I wanted to solve for hydrogen ion concentration and they gave me the pH, well, knowing my logarithmic relationship, it's a base 10. So it's going to be 10 to the negative pH. So if I wanted to solve for hydrogen ion concentration, that formula would be H plus concentration is equal to 10 to the negative pH. If I wanted to solve for hydroxide ion concentration, all I would do is I'd take 10 to the negative pOH. So I can move all the way around this pH scale, right? So again, how many formulas have we been able to come up with here? We came up with H plus equals 10 to the negative pH, pH plus pOH equals 14, pH equals negative log of the hydronium ion, pOH equals negative log of the hydroxide ion, and hydroxide is equal to the... 10 to the negative pOH. One more thing I want to talk about here. Okay, water has the ability to act as what's called amphoteric, meaning it can act like an acid and a base. And there's other molecules that can do that. So when water actually disassociates, right? So let's say here's water. Water disassociates. Let's say it reacts with two waters react. And when it, or we can say it auto ionizes. When it auto ionizes, let's actually just do that, say it's auto ionizing. Alright? When it auto ionizes, it produces an H plus and a hydroxide ion. Okay, that's cool, right? Because I just kind of like, I know what H plus is, and I know what hydroxide is. And I know what water is. We actually were able to calculate a rate at which this actually occurs. The rate at which water auto-ionizes. And we said it's defined as KW. And Kw is equal to 1 times 10 to the negative 14. So that's that rate. And that takes a long time. So that's a very, very small, small number. right? So very very very small number. So with this what's happening is the auto ionization of water you have a specific rate at which this reaction occurs to produce H plus and hydroxide ions. So we can actually say one more thing. Okay well I know that Kw because I'm producing these guys is equal to the H plus concentration times the hydroxide ion concentration. So at any point in time here, I could rearrange this formula. I can get three other formulas from this. I could have this formula, and I could have H plus equals KW over hydroxide. Or I could have hydroxide equals KW over H plus. So there's a lot of different formulas I could derive from that. But this gives me the relationship of KW. All right. Now let's talk about one other thing here, and that's going to be talking about... talking about weak acids and weak bases very briefly, and then we're going to go into a lot of calculations. All right, so we finished up talking about the pH scale. We talked about POH, pH, H+, hydroxide, KW. So what have we finished so far? We've talked about the pH scale. Okay, again, we talked about the logarithmic scale, H+, hydroxide. We talked about acidity, basicity. We talked about KW, right? So we talked about KW, now we got to talk about Ka and Kb, Pk and Pkb. So right now we've only been talking about strong acids and strong bases. So we have to get just a baseline definition of what's up with these weak acids and weak bases. So I'm going to give you an example. First off, let's just do a weak acid. So an example of a weak acid could be, I could pick acetic acid. So I'm going to pick acetic acid. I'm going to say H, C2, H3, O2. Acetic acid is a weak acid. Why is it a weak acid? What was a strong acid? That means that the strong acids are defined as completely disassociating giving all of their protons up. So if I have put in a 0.1 molar solution of hydrochloric acid, all of that 0.1 molar will go over to forming H plus and Cl. If I put a 0.1 molar solution of acetic acid into a solution, not all of it, not all that 0.1 molar will get converted over into H plus and acetate. Why? Because not all of it disassociates. So now what we say is, let's say that we take this reaction, it's a double aryode reaction. So in other words, this can occur, it can actually be in equilibrium. So whenever he gives up his H+, he'll give up his H+, and he gives up and he forms these two structures, and again, what would this be? Weak acid, this would be its conjugate. base. So that's the weak acid conjugate base pair and if I did this in water, water would be the base and H plus would be the conjugate acid. Alright so now because he didn't disassociate completely we have to have some type of number or rate if you will to determine how much of that H plus disassociates. So what do we use for that? That's where Ka comes into play. So Ka is defined as the acid ionization Constant. In other words, it gives us how much of this acetic acid disassociates into a solution. Okay, how much of it goes into H+. Now, here's what we need to know out of this. The greater the Ka, all right, so the more that Ka, the larger it is. The larger the Ka, the stronger the acid. Okay, in other words the bigger this is the more H pluses are going to disassociate. Okay, so if I were to do something like this a low Ka, a low Ka implies a weaker acid. Okay, cool. While I'm here, let's just talk about pKa. So pKa, you can actually calculate it because pKa is the negative log. the Ka. Okay, so now if I said that a Ka, if it was really really big and I wanted to try to be able to figure out pKa, pKa is the actual opposite. So now we say the lower The pKa, the stronger the acid. The higher the Ka, the lower the pKa, the stronger the acid. Because it's going to give you this log amount. And again, we can actually solve for Ka. How? 10 to the negative pKa. So Ka is equal to 10 to the negative pKa. So you can solve for that relationship. So again, higher the Ka, stronger the acid. Lower the pKa, stronger the acid. Lower the Ka, the weaker the acid. The higher the pKa, the weaker the acid. Okay? Now, and again, weaker acid would imply pH. So you could say this would be a low pH, this would be a higher pH. Now let's go ahead and talk about weak bases. So if I talk about weak bases here, let's say I use ammonia. So I use ammonia, so NH3. And I take NH3 and I react him with water. What's the definition of a base? A base accepts protons. So in this case he has to accept a proton from who? Water. So water gives up a proton and puts it onto ammonia and converts him into ammonium which is a weak acid. So there's our base conjugate acid pair. And then H2O gives up a hydrogen so it turns into hydroxide so there is your acid conjugate base pair. Now, same thing, ammonia is not a strong base, so it doesn't accept protons as readily, and it doesn't produce hydroxides as readily as something like sodium hydroxide. So we have to get another type of constant, if you will, to determine how much of this hydroxide is being produced, or how much hydrogens is he accepting. So, we give that a Kb, okay? And Kb is the base. Ionization constant. Okay, and it's the same thing we did up here with the Ka. So the higher the Kb, but remember we got to be careful now, the higher the Kb the stronger the base. Okay, stronger the base. The lower the Kb, the weaker the base. Okay, and the same thing is applied with the PKBs. So PKB here, if I were to say for this one, if it has a high KB, it has to have a low PKB. And if it has a low KB, it has to have a high PKB. All righty. So again, these are your inverse relationships here. So the higher the KB, again, lower the PKB. So that's going to be a stronger base in this condition. And if there's a low KB, it's a weaker base. And if it's a weaker base, it's going to have a high PKB. All right, so we finished off talking about KA and KB now, right? So let's go ahead and say one other thing here before we go into PKAs and PKBs. So I talked about KA and I talked about KB, right? And again, the higher the KA, the stronger the acid, the higher the Kb, the stronger the base. Higher the Ka, the lower the Pk, the higher the Kb, the lower the Pkb, right? And again, that's just a measure of acidity or basicity. Now what we can do is we can say Kw, this is actually working perfectly, Kw is equal to Ka times Kb. So now, why is that? Because if you remember from before, what did we say Kw was equal to? If you remember from before, we said Kw was equal to Kb. We said Kw, which was 1.0 times 10 to the minus 14, is equal to the H plus concentration times the hydroxide ion concentration, right? We talked only about that for strong acids and strong bases. The formula still applies here. But now we just have to take into consideration the amount of the acid that ionizes and then Kb the amount that the base ionizes. So then for this case we can actually say H plus we could apply here the Ka so KW equals Ka and then hydroxide we can say how much of the hydroxide disassociates times Kb. So we could use this formula once again to be able to determine some of these calculations that we're going to do. Okay, so again, KW equals K times KB is an important one. Now let's finish up with PK and PKB. We already know the relationship, right, that PK is equal to the negative log of the KA, and PKB is equal to the negative log of the KB. We can actually say PKA plus PKB equals 14. Doesn't that sound familiar, like pH plus POH equals 14? Same concept here. Here's one thing I want to finish up with with PK and PKBs. We're going to talk about buffers in our final videos. And a buffer is basically a mixture. So a buffer is designed to resist any abrupt changes in the pH. So a buffer is made up of an acid, a weaker acid, so a weak acid, and its associated conjugate base. That's what a buffer is. So when we're picking a buffer. For maybe some type of reaction in the chemistry lab or within the actual human body, we want to make sure that the buffer that we're applying is pretty darn equal to the pH that we desire. So, for example, my blood is naturally going to keep a pH of about 7.35 to 7.45. So, let's say that's the pH. Let's just say 7.4 to keep it right in the middle, right? Keep it a nice number here, 7.4. If I had to pick between a whole bunch of different buffer systems that I'd want in my bloodstream to keep it at that part right there, I'd want the pKa of a buffer to be as close to 7.4 as possible. But we even give it a nice little cushion there. So we say the best buffer is plus or minus one of the pKa, of the pH. So anything that is plus or minus one of the pH is a good buffer. So for example, if I had two buffers here, one of them had a pKa that is equal to 9.3. The other one had a pKa that's equal to, let's say, 6.5. And I had another pKa that's equal to 3.2. Which one of these am I going to pick? I'm going to pick the one that's closest to the pH because that's going to give me the best buffer. Which one's plus or minus one? This guy right there, he's our best choice. So he's going to be the best buffer in this case. Okay, so again, what have we talked about? We talked a lot, a lot of stuff in this video, guys. We talked about the definition of what an acid and a base is according to three different theories. We talked about all of our strong acids and strong bases and what that means to be a strong acid and strong base. We discussed conjugate acid base pairs. We talked about the pH scale and its relationships and formulas. We talked about KW, autoionization of water, and KA and KB. and their relationship to the pH scale. And then we talked about pKa and pKb with respect to buffers and also with respect to Ka. So I hope you guys learned a lot out of this video. It's going to help you guys when we start doing some practice problems. All right, guys, in the next video, we're going to put what we did into action. So the next video, we're going to talk about specifically being able to calculate the pH of specifically strong acids and strong bases. So if you guys want to, click up in the right corner there, and I'll see you guys there.

Before we dive in, let's go ahead and get some just baseline introduction, introductory material out of the way. So first off, let's go ahead and define what an acid and what a base is. Alright, so we're going to do that, then we're going to talk about all the strong acids and strong base. We're going to talk about conjugate acid, conjugate base, pH scale, and we're going to talk about this KWKAKB, PKPKB. Alright, so let's go ahead and dive in first with definition of an acid and a base.

In order for us to do this though, we have to go through three different types of definitions. Alright, so the first one we're going to talk about is Bronsted-Lowry. Okay, so Bronsted-Lowry.

So, Bronsted-Lowry acid. When we're talking about a Bronsted-Lowry acid, he said that when you take an acid and you put it into a solution, an acid gives up protons. So, how does he define an acid?

He defines an acid as a proton donator. Or we can even draw this like this. He loves to give up H plus ions because we're gonna see this a lot, okay, whenever we're doing a lot of these practice problems.

So again, what is a Bronsted-Lowry acid by definition? A proton donator. Well, if an acid is a proton donator, someone has to be able to gain that or accept that proton.

Who accepts the proton? The Bronsted-Lowry base. So the Bronsted-Lowry base is going to be the guy who is accepting the proton. So he's a proton acceptor.

Alright, so that's pretty easy right? So it's not too bad there. So for example if I were to ask you, if I were to draw right here, let's say I did this, H A plus water.

and it yields two products, let's say one of them is A negative and the other one is H3O+. If we were to go based on this definition of Bronsted-Lowry, what happened to this HA? It turned into A negative. So what did it do?

It gave up a proton. So he must be, by definition, the Bronsted-Lowry acid. And then what happened to this H2O? He turned into H3O+, hydronium. So he must have accepted that proton.

And if he accepted the proton that the HA gave, he must be the base. Now while we're here, let's just go ahead and get the... terms out of the way since we're already here conjugate acid conjugate base let's go ahead and get those terms out of the way because we're already here so an acid when he gives up the proton so he donated that proton to the H2O he turns into a different term now that we can call So he was an acid, but now he's a base.

By definition, if he's a base, usually bases have negative charges, right? So they're really good at having negative charges. That's kind of the way that we can see bases is they're negatively charged molecules or atoms.

So we're going to say that this is a base, but it's another type of base. It's the conjugate base. So this acid, it gets converted into a negative, which we're going to call the conjugate.

Base. And again, why is he the conjugate base? Because the acid gives up a proton and turns into a conjugate base. Well, if he turns into the conjugate base, then this must be the acid. Why is he the acid?

Because this water, when he accepts that proton because he's the base, he turns into an acid. So he is the conjugate acid. So if we were to combine the pairs...

This acid and this conjugate base, this is the acid conjugate base pair right there. And this right here is the base conjugate acid pair here. Okay?

So not too bad there, right? So that's just giving you an understanding of Bronsted-Lowry, what an acid and what a base is. And then what is an acid, what is its conjugate base, what is a base, and what is its conjugate acid. Alright, let's go into the next individual who came up with this, another theory. This guy is...

Lewis. So Lewis came up with a theory of acids and bases, but he looked at it in a different way. Bronsted looked at protons, Lewis looked at electrons.

So Lewis said if an acid is donating a proton it has to be accepting electrons. That's his definition. He says by definition a Lewis acid is a molecule who does what? Eps electrons, so he is an electron acceptor.

Okay? So while this guy is donating protons, he's also accepting electrons. Okay cool.

So then a base, what was a base doing? A base was accepting a proton, so he must be donating electrons. So he is an electron donator.

Alright. Not too bad here then, right? So this is, if we were to go based off this, we could apply the same concept here with the Lewis acid and Lewis base. There is two exceptions we're not going to talk about. So Arrhenius, he was the other guy that came up with this theory here, right?

And he said... That an acid is a little bit more specific than just a proton donator. So Arrhenius said an acid is actually going to be a guy who gives up, this guy has to give up a proton, so he is a proton.

proton donator, okay, so he does give up these actual protons, but more specifically, he gives up these protons in an aqueous solution, so we'll say in water. in an aqueous solution. So he even kind of expanded more on that. He said an Arrhenius acid is a proton donator in an aqueous solution.

Okay, the definition of an Arrhenius base, right, so an acid is the one who's going to be able to give up a proton produce an H plus, you know, he's producing a proton in an aqueous solution. A base is just producing hydroxides. So what is it doing? It's producing, or we could say even donating, hydroxides. donator if you'd want to say that or it's producing hydroxides in an aqueous solution.

Okay so that's the definition of an Arrhenius acid and base. An acid is going to be one who gives up protons in an aqueous solution whereas a base is going to be producing hydroxides in an aqueous solution or donating those hydroxides in an aqueous solution. Okay so now we have an understanding of what Bronsted-Lowry acid, Lewis acid and base, Arrhenius acid and base and we even know what conjugate acid and conjugate base pairs are So let's go ahead and check those off.

So what do we do? We did the definition of an acid and a base, and we did what conjugate acids and conjugate bases are. All right. Let's talk about what strong acids and strong bases are. All right.

Well, first off, how would you define a strong acid? A strong acid is defined as something that can deprotonate or disassociate completely. So, for example, I'm going to write down all of our strong acids here.

So let's go ahead and put them right here. here strong acids these guys are going to be hydrochloric acid, hydroiodic acid, hydrobromic acid, HNO3 which is nitric acid, HCl4 which is perchloric acid. H2SO4 which is sulfuric acid.

So how many do we have? 1, 2, 3, 4, 5, 6. These are our six strong acids. And again, what does it mean that they're a strong acid? That means if you guys, if you put this into an aqueous solution, they're going to disassociate so it's a one arrow movement.

If you notice I did not have an equilibrium arrow. Where did I have an equilibrium arrow? Here. Here I will not because all of the HCl will deprotonate into the form of H plus and Cl negative.

And again, a strong acid is defined as a proton donator. So it's going to give up that proton and it completely disassociates in an aqueous solution. So what would this one look like?

HI would give way to H plus and I negative. HBr would give way to H plus and Br negative. HNO3 would give way to H plus. and nitrate. HCl4 would give way to H plus and ClO4 negative.

H2SO4 is different. So if you look here, H2SO4, it has an H2. This is defined as a diprotic.

acid, meaning he can give up two protons. So if we were to run this reaction, he would deprotonate and we'd turn into HSO4 negative. And then he would react even further and get converted into SO4 2 negative, right?

So what we give up, we should actually show the H plus in there. So again, it would react here and give you H plus and this guy. Then this would react even further and produce sulfate and H plus. So again, these are six strong acids and the reason why they're strong acids is because they disassociate completely.

Alright, let's do strong bases. bases same thing the reason they're called strong bases is because they disassociate completely when put in an aqueous solution so what are this what are these strong bases these ones are easy because if you take all the group one and group two metals there is a couple exceptions but if you take almost all of those so for example we take sodium hydroxide we take lithium hydroxide we take potassium hydroxide we take rubidium hydroxide We can even keep going, we can even do cesium, okay, cesium hydroxide. We can't do francium though. Alright, that's a big mofo and not a good one to do, right?

Alright, now what's another one? We can do group 2 metals. So what are the group 2 metals?

The group 2 metals are the ones that are in the middle. metals we could do magnesium isn't red necessarily soluble calcium is too but it's kind of slightly soluble but we'll put that one in there COH 2 we could also do strontium hydroxide and we can could even do barium hydroxide but not radium. Okay so in general these are our strong bases.

Magnesium isn't really considered to be a strong base because it's not really soluble. Okay so if we were to kind of look at all of these guys here this would pretty much give us all of our strong bases and what are they strong why the strong bases because they can give up or donate all of these hydroxides all of them will disassociate into solution. Okay so if you were to put all of these guys into a solution I'm only gonna do one here if I were to take sodium hydroxide here, put him into an aqueous solution, what would happen?

It would give up or completely disassociate into hydroxide ions. Okay, so that's the definition of a strong base and that's the definition of a strong acid. They completely disassociate. It's not a reversible reaction. One more relationship I want to develop here because this is really important when we're going through a lot of these calculations.

There is a relationship here and it says the stronger The acid, whenever the reaction occurs, the weaker its conjugate base. The weaker its conjugate base. And this would be the vice versa if you start off with a strong base, the weaker its conjugate acid.

It's either way. Let me give you an example here. If I take something like HCl, I react this in water, what would I get?

I would get H+, which is a strong acid. acid and I'll get chloride. Well this by definition is a strong acid. This right here is a very very weak conjugate base.

That's going to be important whenever we talk about the acidity of and basicity of salts. Okay, alright so again remember that relationship stronger the acid the weaker its conjugate base and like I said you could do the same thing with the other one. You could say stronger the base The weaker its conjugate acid. Okay, so now we have that relationship taken care of.

So what have we done so far? We've talked about definitions of acids and bases with Bronsted, Lewis, and Arrhenius. We talked about all the strong acids and strong bases, and we tied in the relationship with conjugate acid, conjugate base. and the relationship of strength. Alright let's get over here to the pH scale.

Alright so let's go ahead and talk about pH scale. So pH scale is a is basically a logarithmic scale here and we're going to talk about that a little bit but just know that that it's a scale that goes from 0 all the way to 14 and what it does it helps to measure from 0 all the way to 7 measures how acidic the solution is and then going from 7 all the way up to 14 measures how basic a solution is now again what do I mean by acidic and what do I mean by basic So when we talk about acidic, we're talking about specifically, we're talking about how many protons we have or how many hydronium ions we have. So when we talk about acidic environments, we're talking about the concentration of...

of H plus and we put these brackets in there to signify concentration. So the more you go closer to zero, the more acidic your solution is. So if you have a solution that's in between one and zero, that's pretty freaking acidic.

So when we talk about acidity, we're talking about H plus concentration. When we're talking about basicity, we're talking about Hydroxide ion concentration. So again, if I were to say, okay, I got a solution that can have H pluses and I got a solution that can have hydroxides. How do I know specifically that it's acidic? Well, that means that the H plus concentration in the acidic environment is greater than the concentration of the hydroxide ion.

Alright, specifically in the acidic environment because the more H plus is, the more acidic. In the basic solution, the basic solution we want it to be measured by hydroxide ion concentration. So the greater the hydroxide ion concentration is, then the H plus concentration, that means that we're in a basic solution.

What if it's neutral though? Then these two are equal. So then the H plus concentration is equal to the Hydroxide ion concentration and this is an important relationship when we talk about Kw here in a second.

Okay before we do that, okay I want to talk about another thing. So I said before that this is a logarithmic scale right. So another words if I wanted to go from zero all the way to five most people would say okay how many people was the pH change well they'd be like oh it was five that's wrong and the reason why is each number you go up from zero to one 1 to 2, 3 to 4, and then 4 to 5 is a 10 times change.

It's a 10-fold change. So, for example, if I'm going up here, I'm actually going up what? 10 times Okay, so what is that? 10 times 10 times 10 times 10 times 10. Okay, well that's going to be what? That's going to be a 100,000 fold change.

That is a 100,000 fold change. So how much of a change was it? It was a 100,000 times change. So remember that whenever we're saying okay, what's the pH change going from 0 to 5?

Well, each change is a tenfold increase So if I go 10 and 10 and 10 and 10 and 10 that is a five-fold increase But specifically by 10 times 10 times 10 times 10, right? That's just a hundred thousand times change there same thing if I was going from 14 to 9 If I'm going all the way from 14 to 9, what am I doing? It's changing by a 10-fold scale. So 10 times 10 times 10 times 10, and then over here to this guy times 10. So same situation, how many was that? We went from 14 all the way to 9, most people would be like, oh, that's a 5 times difference.

No, it's a 100,000 times difference, because it's a 10 times each one, okay? Now another thing, because of this base 10, if you will, change, they came up with a way to be able to calculate a specific number which we call this pH, right? And so what they said is, okay, well, we can technically figure out what the pH of a solution is since we have this scale now.

And what they said is we could actually say that pH is equal to... We put a negative in there for specifically mathematical sake, and we said it's a log base 10 relationship. It's a logarithmic scale.

So pH is equal to the negative log of the H plus concentration. Okay, and that's how we derive this formula to solve for pH. We said it's based on this logarithmic scale, which is a base 10, right? Tenfold change.

Same thing, that's pH. So that would give you the H plus concentration. P-O-H. would give you the hydroxide ion concentration.

So it's the negative log of the hydroxide ion concentration. Okay, well that's cool because now I have this ability to calculate pOH, I have the ability to calculate pH. What if I add both of these two together? So then I take and I add pOH plus pH. That gives me a specific formula. pH plus pOH, wait, isn't that the whole scale? It is.

So it equals 14. And that's how we got a lot of these formulas, right? So what do we have so far? We have pOH equals negative log of the hydroxide ion, pH equals negative log of the hydronium ion, and then pH plus pOH equals 14. In the same way, if you know your algebra, if I wanted to solve for hydrogen ion concentration and they gave me the pH, well, knowing my logarithmic relationship, it's a base 10. So it's going to be 10 to the negative pH. So if I wanted to solve for hydrogen ion concentration, that formula would be H plus concentration is equal to 10 to the negative pH. If I wanted to solve for hydroxide ion concentration, all I would do is I'd take 10 to the negative pOH. So I can move all the way around this pH scale, right?

So again, how many formulas have we been able to come up with here? We came up with H plus equals 10 to the negative pH, pH plus pOH equals 14, pH equals negative log of the hydronium ion, pOH equals negative log of the hydroxide ion, and hydroxide is equal to the... 10 to the negative pOH.

One more thing I want to talk about here. Okay, water has the ability to act as what's called amphoteric, meaning it can act like an acid and a base. And there's other molecules that can do that.

So when water actually disassociates, right? So let's say here's water. Water disassociates. Let's say it reacts with two waters react.

And when it, or we can say it auto ionizes. When it auto ionizes, let's actually just do that, say it's auto ionizing. Alright? When it auto ionizes, it produces an H plus and a hydroxide ion. Okay, that's cool, right?

Because I just kind of like, I know what H plus is, and I know what hydroxide is. And I know what water is. We actually were able to calculate a rate at which this actually occurs. The rate at which water auto-ionizes. And we said it's defined as KW.

And Kw is equal to 1 times 10 to the negative 14. So that's that rate. And that takes a long time. So that's a very, very small, small number.

right? So very very very small number. So with this what's happening is the auto ionization of water you have a specific rate at which this reaction occurs to produce H plus and hydroxide ions. So we can actually say one more thing.

Okay well I know that Kw because I'm producing these guys is equal to the H plus concentration times the hydroxide ion concentration. So at any point in time here, I could rearrange this formula. I can get three other formulas from this. I could have this formula, and I could have H plus equals KW over hydroxide.

Or I could have hydroxide equals KW over H plus. So there's a lot of different formulas I could derive from that. But this gives me the relationship of KW. All right.

Now let's talk about one other thing here, and that's going to be talking about... talking about weak acids and weak bases very briefly, and then we're going to go into a lot of calculations. All right, so we finished up talking about the pH scale.

We talked about POH, pH, H+, hydroxide, KW. So what have we finished so far? We've talked about the pH scale. Okay, again, we talked about the logarithmic scale, H+, hydroxide.

We talked about acidity, basicity. We talked about KW, right? So we talked about KW, now we got to talk about Ka and Kb, Pk and Pkb. So right now we've only been talking about strong acids and strong bases.

So we have to get just a baseline definition of what's up with these weak acids and weak bases. So I'm going to give you an example. First off, let's just do a weak acid.

So an example of a weak acid could be, I could pick acetic acid. So I'm going to pick acetic acid. I'm going to say H, C2, H3, O2. Acetic acid is a weak acid. Why is it a weak acid?

What was a strong acid? That means that the strong acids are defined as completely disassociating giving all of their protons up. So if I have put in a 0.1 molar solution of hydrochloric acid, all of that 0.1 molar will go over to forming H plus and Cl. If I put a 0.1 molar solution of acetic acid into a solution, not all of it, not all that 0.1 molar will get converted over into H plus and acetate. Why?

Because not all of it disassociates. So now what we say is, let's say that we take this reaction, it's a double aryode reaction. So in other words, this can occur, it can actually be in equilibrium. So whenever he gives up his H+, he'll give up his H+, and he gives up and he forms these two structures, and again, what would this be?

Weak acid, this would be its conjugate. base. So that's the weak acid conjugate base pair and if I did this in water, water would be the base and H plus would be the conjugate acid.

Alright so now because he didn't disassociate completely we have to have some type of number or rate if you will to determine how much of that H plus disassociates. So what do we use for that? That's where Ka comes into play.

So Ka is defined as the acid ionization Constant. In other words, it gives us how much of this acetic acid disassociates into a solution. Okay, how much of it goes into H+. Now, here's what we need to know out of this. The greater the Ka, all right, so the more that Ka, the larger it is.

The larger the Ka, the stronger the acid. Okay, in other words the bigger this is the more H pluses are going to disassociate. Okay, so if I were to do something like this a low Ka, a low Ka implies a weaker acid.

Okay, cool. While I'm here, let's just talk about pKa. So pKa, you can actually calculate it because pKa is the negative log. the Ka. Okay, so now if I said that a Ka, if it was really really big and I wanted to try to be able to figure out pKa, pKa is the actual opposite.

So now we say the lower The pKa, the stronger the acid. The higher the Ka, the lower the pKa, the stronger the acid. Because it's going to give you this log amount.

And again, we can actually solve for Ka. How? 10 to the negative pKa. So Ka is equal to 10 to the negative pKa. So you can solve for that relationship.

So again, higher the Ka, stronger the acid. Lower the pKa, stronger the acid. Lower the Ka, the weaker the acid.

The higher the pKa, the weaker the acid. Okay? Now, and again, weaker acid would imply pH. So you could say this would be a low pH, this would be a higher pH. Now let's go ahead and talk about weak bases. So if I talk about weak bases here, let's say I use ammonia.

So I use ammonia, so NH3. And I take NH3 and I react him with water. What's the definition of a base? A base accepts protons.

So in this case he has to accept a proton from who? Water. So water gives up a proton and puts it onto ammonia and converts him into ammonium which is a weak acid.

So there's our base conjugate acid pair. And then H2O gives up a hydrogen so it turns into hydroxide so there is your acid conjugate base pair. Now, same thing, ammonia is not a strong base, so it doesn't accept protons as readily, and it doesn't produce hydroxides as readily as something like sodium hydroxide.

So we have to get another type of constant, if you will, to determine how much of this hydroxide is being produced, or how much hydrogens is he accepting. So, we give that a Kb, okay? And Kb is the base.

Ionization constant. Okay, and it's the same thing we did up here with the Ka. So the higher the Kb, but remember we got to be careful now, the higher the Kb the stronger the base. Okay, stronger the base. The lower the Kb, the weaker the base.

Okay, and the same thing is applied with the PKBs. So PKB here, if I were to say for this one, if it has a high KB, it has to have a low PKB. And if it has a low KB, it has to have a high PKB.

All righty. So again, these are your inverse relationships here. So the higher the KB, again, lower the PKB.

So that's going to be a stronger base in this condition. And if there's a low KB, it's a weaker base. And if it's a weaker base, it's going to have a high PKB.

All right, so we finished off talking about KA and KB now, right? So let's go ahead and say one other thing here before we go into PKAs and PKBs. So I talked about KA and I talked about KB, right? And again, the higher the KA, the stronger the acid, the higher the Kb, the stronger the base.

Higher the Ka, the lower the Pk, the higher the Kb, the lower the Pkb, right? And again, that's just a measure of acidity or basicity. Now what we can do is we can say Kw, this is actually working perfectly, Kw is equal to Ka times Kb.

So now, why is that? Because if you remember from before, what did we say Kw was equal to? If you remember from before, we said Kw was equal to Kb.

We said Kw, which was 1.0 times 10 to the minus 14, is equal to the H plus concentration times the hydroxide ion concentration, right? We talked only about that for strong acids and strong bases. The formula still applies here. But now we just have to take into consideration the amount of the acid that ionizes and then Kb the amount that the base ionizes.

So then for this case we can actually say H plus we could apply here the Ka so KW equals Ka and then hydroxide we can say how much of the hydroxide disassociates times Kb. So we could use this formula once again to be able to determine some of these calculations that we're going to do. Okay, so again, KW equals K times KB is an important one. Now let's finish up with PK and PKB. We already know the relationship, right, that PK is equal to the negative log of the KA, and PKB is equal to the negative log of the KB.

We can actually say PKA plus PKB equals 14. Doesn't that sound familiar, like pH plus POH equals 14? Same concept here. Here's one thing I want to finish up with with PK and PKBs.

We're going to talk about buffers in our final videos. And a buffer is basically a mixture. So a buffer is designed to resist any abrupt changes in the pH. So a buffer is made up of an acid, a weaker acid, so a weak acid, and its associated conjugate base. That's what a buffer is.

So when we're picking a buffer. For maybe some type of reaction in the chemistry lab or within the actual human body, we want to make sure that the buffer that we're applying is pretty darn equal to the pH that we desire. So, for example, my blood is naturally going to keep a pH of about 7.35 to 7.45. So, let's say that's the pH. Let's just say 7.4 to keep it right in the middle, right? Keep it a nice number here, 7.4.

If I had to pick between a whole bunch of different buffer systems that I'd want in my bloodstream to keep it at that part right there, I'd want the pKa of a buffer to be as close to 7.4 as possible. But we even give it a nice little cushion there. So we say the best buffer is plus or minus one of the pKa, of the pH. So anything that is plus or minus one of the pH is a good buffer.

So for example, if I had two buffers here, one of them had a pKa that is equal to 9.3. The other one had a pKa that's equal to, let's say, 6.5. And I had another pKa that's equal to 3.2.

Which one of these am I going to pick? I'm going to pick the one that's closest to the pH because that's going to give me the best buffer. Which one's plus or minus one?

This guy right there, he's our best choice. So he's going to be the best buffer in this case. Okay, so again, what have we talked about?

We talked a lot, a lot of stuff in this video, guys. We talked about the definition of what an acid and a base is according to three different theories. We talked about all of our strong acids and strong bases and what that means to be a strong acid and strong base. We discussed conjugate acid base pairs. We talked about the pH scale and its relationships and formulas.

We talked about KW, autoionization of water, and KA and KB. and their relationship to the pH scale. And then we talked about pKa and pKb with respect to buffers and also with respect to Ka. So I hope you guys learned a lot out of this video. It's going to help you guys when we start doing some practice problems.

All right, guys, in the next video, we're going to put what we did into action. So the next video, we're going to talk about specifically being able to calculate the pH of specifically strong acids and strong bases. So if you guys want to, click up in the right corner there, and I'll see you guys there.

Transcript for:Understanding Acids and Bases in Chemistry

Transcript for:
Understanding Acids and Bases in Chemistry