Understanding Acids and Bases in Chemistry

An introduction to acids and bases is going to be the topic of this lesson. My name is Chad, and welcome to Chad's Prep, where my goal is to take the stress out of learning science. Now, in addition to high school and college science prep, we also do MCAT, DAT, and OAT prep as well. I'll leave links in the description for you can find those courses. Now, this lesson is part of my new general chemistry playlist. I'm releasing several lessons a week throughout the school year, so if you want to be notified every time I post one, subscribe to the channel, click the bell notification. Okay, in this introduction to acids and bases, we're going to... Define acids and bases. Well, it turns out we're going to look at three different definitions, and we're going to use one of those definitions to talk about what are called conjugate acids and conjugate bases. And then finally, we'll polish this lesson off by distinguishing between strong acids and bases and weak acids and bases. So let's start off with those three definitions. So the three definitions for acids and bases we've got are Arrhenius, Bronsted-Lowry, and Lewis. And we're going to kind of go in historical order here, as well as this kind of order of progression where they get to be a more complete definition. And what I mean by this is that Not all acids and bases will actually satisfy Arrhenius'definition. He would not have considered all acids or bases that we now know of to be acids or bases. Now, Bronson Lowry's definition is a little broader and includes a greater number of acids and bases, but Lewis has the ultimate definition. Every acid or base that we currently, in modern times, identify as acids or bases, Lewis'definition will include all of them, but the other two, not quite so much. Let's talk about this for a second. So we start with Arrhenius here. So Arrhenius first off, he only dealt with aqueous solutions. That's it. And so everything had to be in water for Mr. Arrhenius here. And the way he looked at an acid is an acid is anything that causes the H3O plus concentration to go up in water. That's what he looked at. So let's take a look at what that ultimately means here. So Most famous acid. I'm gonna ask you, what is the most famous acid? And you're not supposed to say LSD. It is HCl, hydrochloric acid, found in lots of chemistry labs, so, but also found in your stomach. It's the same stuff that you put in your pool that might also be called muriatic acid. So it is the most famous acid, not LSD. In this case, if you look at the reaction of what happens when you put HCl in water, so... HCl donates one of its H plus ions over to water to form H3O plus. So we call H3O plus hydronium. And so it turns out that anything that when you place it in water causes you to form hydronium, that's what Arrhenius looked at as an acid here. And so in this case again, increases the concentration of hydronium. You might also consider this me, you know, something kind of like an H plus donor in water. But the key is it had to be in water for Arrhenius to be considering this. Now if we look at bases on the other hand, So bases instead increase the concentration of hydroxide. Instead, the most famous base is NaOH, sodium hydroxide, and it simply dissociates to produce OH-hydroxide ions. Now it turns out when you've got an aqueous solution, so usually you have an equal concentration of H3O+, and OH-. We'll study this a little more in subsequent lessons in this chapter. So we've got equal amounts. So, but it turns out if you add an acid, you're going to end up with more H3O plus than OH minus. And if you add a base, then you're going to end up with more OH minus than H3O plus. And there's this interplay between the two. So, but that's what Mr. Ranius considered here. Whereas Mr. Bronsted-Lowry came along and said, hey, we've got a little bit better definition. So for one, this doesn't have to only be about water. We can use any solvent. We can use alcohols and things of this sort. In fact, we don't even have to use a solvent. We can do acid-base reactions in the gas phase. So, and the way they looked at this then, is that an acid was simply an H plus donor. Didn't have to be in water at all. It's just an H plus donor. And another word for H plus is a proton. So, proton donor, if you notice that hydrogen has one proton in its nucleus and then one electron. And if you lose the electrons that you become H plus, then all you are is that proton in the nucleus. And so that's why H plus is often referred to not simply as the hydrogen ion, but also just simply as a proton. And so another way to say this is proton donor here for a Bronsted-Lowry acid. That's the way they looked at it. On the base side, they kind of said, well, you know, rather than looking, you know, being about hydroxide, which would be, you know, kind of all about aqueous solutions, it's really just about this transfer of this H plus ion of this proton. And so, whereas the acid is the proton donor, the base is the proton acceptor. And so we're really just looking at the transfer of protons. So if we go back to this, reaction with aqueous solutions. So Bronsted and Lowry would look at this and say, okay, between the two reactants, who gave away the H plus? Well, that would be the HCl. So they define him as the acid. And who accepted the H+, who gained the H+, well, H2O turns into H3O+, gaining not just another hydrogen, but a plus charge at the same time, and that makes him the base. And so, again, this is happening in water, and so it's, you know, satisfying Arrhenius'definition of an acid-base reaction, stuff like that. So, however, it's also satisfying Bronsted-Lowry's. However, again, Bronsted-Lowry says, you know, this, you know, The whole idea of acid bases doesn't have to involve water. And so in fact, we could mix HCl and ammonia in the gas phase. These could be gases, not aqueous solutions at all. And you're still going to get a reaction going on here. HCl is still the acid. And in this case, ammonia would be the base. And so we take a look at this and who's the proton donor? Well, again, that's going to be the HCl. He's giving away an H plus ion to become just plain old Cl minus. And who's accepting that or receiving it? NH3 turning into NH4 plus in this case, so he is the base. And so I just want you to see that all the acid-base reactions that take place in water that Arrhenius would have called acid bases, Bronsted and Lowry will also call acid and bases. However, even the reactions that don't involve water, so Bronsted and Lowry are now gonna consider those acid-base reactions as well. And again, it's just the idea that we're just transferring a proton, follow that proton. So who gives away the H plus and who receives or accepts the H plus is how we'll identify acid base. And it turns out this is the most useful definition, kind of working definition for identifying acid bases. So it's much easier to follow the H than what we're going to see with Mr. Lewis, his idea in a second. So when I'm having you identify acid bases throughout this chapter, most of the time what I'm expecting you to do is use Bronsted and Lowry's definition to do so. All right, so Mr. Lewis came along and said, hey, you know, your definition is really good, Mr. Bronson and Lowry, but I've got a few acids that you're not gonna classify as acids that I still will. And so if we take a look, he said, it's not about the protons, transfer protons, all about electrons. And let's not put that in blue. And so rather than an acid being a proton donor, it's actually that an acid is an electron acceptor. And rather than a base being a proton acceptor, it is an electron donor. And sometimes you'll hear electron pair acceptor and electron pair donor. And the idea is that we're making a new bond in an acid-base reaction. So, and as long as that new bond is being made to a hydrogen ion, an H plus ion, well then you'd consider that acid-base reaction not only a Lewis acid and base reaction, but a Bronsted-Lowry acid-base reaction, no big deal. But if the new bond being made is to something other than hydrogen, then Bronsted-Lowry's definition won't apply, but Lewis's still will. So let's take a look at this here for a second. Let's take a look at this last one. Let's make some room. Now let's look at this instead from Lewis's perspective. So Mr. Lewis here, it helps if we actually draw some Lewis structures. Those are the three that we're going to need here to take a look at. So, and you want to look and identify where you've made a new bond in the reaction, and that's right there in our case. That's the new bond. And that new bond is between nitrogen and hydrogen. It's between this nitrogen right here and the hydrogen that was donated by HCl, again at least from Bronsted and Lowry's perspective. And so the question is, where did the electrons come from to make that bond? Did they come from nitrogen or did they come from hydrogen? Well, they came from nitrogen. Nitrogen uses its lone pair of electrons. It donates those electrons to be the ones that nitrogen and hydrogen will share. And so in that sense, he's the electron pair donor. And that, according to Mr. Lewis, would make him the base. And HCL says, hey, I didn't contribute anything to make this bond. I just accepted your offer to make me a bond. So I'm the electron pair acceptor. Thank you very much. Cool. That makes him the Lewis acid. And so if you classify anything as an acid or a base, Under any of the less complete definitions, they will still be classified as acids or bases under Mr. Lewis's definition, but sometimes it's a little harder to see. Notice I actually had to draw the Lewis structures to kind of figure out who's the Lewis acid and who's the Lewis base. Now there's also a class of acids that are just going to act as Lewis acids, but they're not going to be a part of any other definition. Arrhenius wouldn't have considered them necessarily acids or Bronsted-Lowry wouldn't consider acids, but Lewis has a small class of acids that he would consider acids. We want to take a look at those as well. In fact, we'll just do that right over here. So if we take a look, boron compounds are really commonly gonna fall into this category. I'm gonna leave off a lot of the lone pairs here that are on the fluorines. We don't need those for this reaction. And so boron, you might recall, is famous for not having a filled octet. And not having a filled octet, boron actually has an empty orbital. And an empty orbital makes you a good candidate for being an electron acceptor. It turns out, a place where you can accept those electrons. And so boron-containing compounds are often going to be able to act as Lewis acids, and only going to be acting as a Lewis acid by any of the definitions. Won't satisfy either the Uranus or Bronsted-Lowry definition. definition, but also metal ions, it turns out. Metal ions, being cations, have lots of empty orbitals in their valence shell quite commonly, and so metal ions will often act as electronic scepters, Lewis acids as well, even though, again, Arrhenius and Bronsted-Lowry would not have necessarily considered them acids. So, just want to give you an idea of, you know, if I ask you to identify an acid that only satisfies Lewis's definition only, so... probably going to be a boron containing compound or a metal ion like aluminum ions are a good example. And so if we look at what actually happens here in this reaction, we end up with a positive formal charge of nitrogen and negative formal charge on boron. And the new bond we made is the bond between nitrogen and boron. And the question is, well, where did those electrons come from? From nitrogen or from boron? Well, they're again coming from the lone pair on the nitrogen. Boron didn't have a lone pair to donate here. And so the hallmark here again of our Lewis acid, it's going to have empty orbitals. And that's usually again going to be, if it's going to be uniquely a Lewis acid, it's going to be a boron containing compound or a metal ion. But we also saw that, you know, HCl satisfied, you know, the qualifications for being a Lewis acid as well. If you're an Arrhenius acid or Bronsted-Lowry acid, then you are also a Lewis acid. Again, Lewis had the ultimate definition. Now, the hallmark of being the Lewis base here, and you should recognize this, is to act as a Lewis base, you need a lone pair of electrons. That is the one qualification you must have. Cool. And as a Lewis base, if you use that lone pair of electrons to make a new bond to a hydrogen, again, then not only are you a Lewis base, you're going to be a Bronsted-Lowry base as well. That's kind of the distinguishing factor. But if you use that lone pair of electrons to make a bond to any other atom besides hydrogen like boron, then you're going to be a Lewis base, but you're not going to be a Bronsted-Lowry base in that case. So those are your three definitions. You should definitely know verbatim what these definitions are, but you probably want to know how to apply them as well. And again, as far as working definition for identifying acids and bases, We're going to use Bronsted-Lowry's definition more commonly than any other. However, you should keep in mind that Lewis has the ultimate definition. If you are an acid, there's a good chance you're probably an Arrhenius acid. There's a really good chance that you're probably a Bronsted-Lowry acid, but there is a 100% chance that you're a Lewis acid. He has the most complete definition. All right, so I left the Bronsted-Lowry definitions for acid bases on the board here because... That's going to be crucial for us identifying what are referred to as conjugate acid conjugate base pairs. And so, it turns out this is going to have some relevance when we start talking about trends in acidity and trends in basicity and stuff like this. And it might just really be a question on the test that you have to identify an acid's conjugate base or a base's conjugate acid. And so it turns out you've often got these two species in equilibrium, a conjugate acid conjugate base pair. in equilibrium together in most solutions and stuff like this. You might have more of one than the other depending on conditions and stuff like this, but they're usually in equilibrium together, and their relationship often determines how strong the acid is or how strong the base is, things of this sort. So super important concept to identify, and it is Bronsted-Lowry's definition that's going to allow us to identify these. So if you take a look at HCl here, I want to fill in this table here, and we want to know what is HCl's conjugate base. So here's the conjugate acid. The question you get on the test might just be, What is the conjugate base of HCl? Well what you want to do here is you just want to have HCl act as an acid. Well what do acids do according to Mr. Bronson-Lowry? Well they're H plus donors. So he's going to donate an H plus and whatever you have left when he's done, that is your conjugate base. Well after he gets rid of the H plus, he has one fewer H's. So now we have a chlorine atom with no H's, only had one to start with, and one fewer plus charge. Well, he was a neutral species before, and so if you lose a positive charge, you're now going to have a negative charge. And so Cl-would be properly identified as the conjugate base of HCl. Cool. And the big thing is that after HCl, which again, most famous strong acid, after he acts as an acid, what he turns into is his conjugate base. Cool, let's go the other way here. So we've got ammonia next on the chart listed as a conjugate base. And the question you might see on a test is which of the following is the conjugate acid of ammonia. So we'll start with ammonia here. And in this case, we just need him. He's been identified as the base here because we're looking for his conjugate acid. He must be the base. And so we just need him to act as a base. Well, what do bases do? Well, they're H plus acceptors. So we just need him to accept. an H+. And so if he accepts an H+, meaning he gains an H+, so instead of having three of those H's, he's now going to have four, but he also gains a positive charge as well. So he was a neutral molecule before, he's now going to have an overall plus one charge. And that NH4 plus would be the conjugate acid here we'd fill in on the chart. And so the question again, if the question was, what is the conjugate acid of NH3? you'd say NH4+. Okay, so now we've got to get a little bit tricky here. And the way we're going to get tricky is by introducing what we're going to call amphiprotic. So you might think of ampha like an amphibian as an animal that can live on land or water. Well, amphiprotic is a species that can act as a proton donor or a proton acceptor. it can act as a Bronsted-Lowry acid or a Bronsted-Lowry base. That's what amphiprotic means. And we're going to choose such compounds because if I asked you to identify a conjugate acid or a conjugate base, well, amphiprotic species have both a conjugate acid and a conjugate base. And so if you don't realize which one you're being asked or get it backwards in your mind, it's a great question to kind of slip you up on the test. And so in this case, we're going to start with bicarbonate here, HCO3-. And so it turns out bicarbonate can act as an acid and donate another H plus ion. or it can act as a base and accept an H plus ion. And so you gotta really know your definitions here so you can figure this out. So if we start with bicarbonate here and we say, what is the conjugate base of bicarbonate? What goes in this box? Well, if I'm asking for the conjugate base, then it can be implied that the bicarbonate itself must be the conjugate acid form. And what do acids do? Acids donate H plus, they lose that H plus. And so if we lose an H plus from the formula here, then we'll have no Hs left. And again, losing a plus charge, well, we're already negative one, so we're going to be one less positive, i.e. one more negative, and so we'll have a negative two charge. And so we get the carbonate ion. That is the conjugate base of the bicarbonate ion. But you could ask this backwards. Again, bicarbonate is amphiprotic. It can act as an acid or base. And so if instead, we want to know what goes in this box. What is the conjugate acid of the bicarbonate ion of HCO3-? Well, again, if I want the conjugate acid, then I got to know that HCO3-is acting as the base. He's the conjugate base this time. What do bases do? Well, they're H plus acceptors. They gain H plus. And so in this case, if HCO3-gains an H, he's now going to have two of them. But he also gains a plus charge again. And so if he started out as minus one, he's now going to be neutral. And H2CO3 there, carbonic acid, is the conjugate acid of HCO3-bicarbonate. Now, if I want to get super tricky on this, I will do the same thing with hydroxide. We normally don't think of hydroxide as being amphiprotic. Hydroxide is pretty much like your classic base. In an aqueous solution, hydroxide is kind of how you recognize you have base in the solution. And so whenever you hear hydroxide, it's going to be really kind of this association in your mind that it's a base. And so if I said, hey, what's the conjugate of hydroxide? You're not even going to think like, oh, you didn't. You missed a word there, Chad. You just said conjugate. You didn't say conjugate acid or base. Which one did you mean? You're just going to think, oh, he must have just meant conjugate acid. Well, and that could be the question. What is the conjugate acid of hydroxide? Well, if I'm looking for the conjugate acid, then hydroxide must be acting as the conjugate base. And what do bases do? They're H plus ion acceptors. And if OH minus accepts an H, you're now going to have two Hs with that oxygen. And it's going to gain a plus charge to now be neutral. And that's water. And so water is the conjugate acid of hydroxide. But if I was trying to be tricky because, you know, maybe my team just lost the Super Bowl and I'm angry at all my students, or at least I'm going to take it out on them. So then I might ask, what's the conjugate base of hydroxide? Well, we don't normally see hydroxide acting as an acid, but if I'm asking for his conjugate base, then he must be the conjugate acid. This is super uncommon. It turns out you can get a hydroxide to act as an acid if you react it with an even stronger base, which is rare and not something you're likely to encounter in this class, but in organic chemistry you might. So, but hydroxide, if he's the acid and I'm looking for his conjugate base, well then have him act as the acid, have him donate another H plus ion. And if he gets rid of an H, then he won't have any left, that oxygen, and he'll also lose a plus charge. So instead of being minus one, he's now going to be minus two and you get the oxide ion. And so the conjugate base of hydroxide is the oxide anion here, O2 minus. Super tricky. I've seen it show up a time or two on an exam. Only reason I'm covering it. Cool. So hopefully now you can identify conjugate acids and conjugate-based pairs. And if I give you one, you can ask for the other and things of this sort. You might also want to be able to do one other thing. Alright, so I'm going to try and give you an example that's not so obvious. And you might have noticed I've been using a lot of red and blue. And red for acids and blue for bases. So I'm not doing that here. So, and if you don't know what this is from, it's from litmus paper. Litmus paper turns red in the presence of acids, and it turns blue in the presence of bases. It's a nice little quick test to see if you've got an acid or a base. Well, here I'm giving this to you all in black, so you can't figure it out. So, and I'm giving you an acid-base reaction that is reversible, and most acid-base reactions, it turns out, are reversible. And so I want you to identify the Bronsted-Lowry acids and bases in this reaction. And you can look at it going in the forward direction, but it is reversible, and you can look at it going in the reverse direction. And so what you're going to find is in a typical acid-base reaction, according to Mr. Bronsted and Lowry here, you're going to have an acid base on the reactant side going in the forward direction, but you'll have an acid base in the product side going in the reverse direction, and you just got to figure out which based on your definitions. And so if we look at this in the forward direction, who is gaining the H plus ion? Well in this case that's NH3. He's going from NH3 to NH4, and we can see that CH3COOH is the one losing the H, and it's that last H plus ion that he's losing. in the process. And so the one gaining it, the H plus acceptor, is our base. And the one losing it is our acid. But notice, identifying the ones on the other side, well, in this case, again, if ammonia here is the base, well, after he acts as a base, what do you turn into? Well, after you act as your base, you turn into your conjugate acid. And so on this side, if I was doing the reverse direction, NH4 plus would be the one donating the H plus. And I see my conjugate base conjugate acid pair right there. So, and same thing the other way, if I was doing the reverse reaction, it'd be CH3COO minus, that would be gaining the H. And he's the conjugate base of this acid right here. Cool, and so in a typical Bronsted-Lowry acid-base reaction, you should be able to identify the acid and base both on the reactant and the product side as well. All right, so the last topic in this lesson, we've got to talk about the difference between strong acids and bases and weak acids and bases. And the difference in strength, it turns out, is how much they dissociate. So if you're a proton donor, a Bronsted-Lowry acid, it's how much you donate that proton. That's what we're going to refer to as being as dissociation or ionization. All right, so if we take a look, we've got seven strong acids on our list. I will let you know that most of you are either going to have six or seven. So, and that last one, HClO3, it shows up on some lists, and it doesn't show up on others. It's kind of like a, right at the cusp of being strong, and some people put it on their list, and some people don't. Now, there are not only seven strong acids in the universe, but there's not too many more, it turns out. But these are the seven you're likely to see, again, if you're not, in some cases, going to be given only six. And it turns out, if you see an acid that's any other than these, you should identify it as a weak acid, which means it will not dissociate 100%. So... strong acids associate 100% in water. So if we go back to our most famous acid again, HCl, we saw its reaction with water earlier to form hydronium and its conjugate base, and this effectively goes 100% to completion. Now truth be told it's not 100%, but it's like 99.99%, so it's effectively 100% to completion. And the truth is that would happen for any of these acids. So if you put any of these acids in water, they disappear completely because they dissociate completely. And the only acid that will actually be present in that solution will be H3O+. And so it turns out when you have aqueous solutions of any of the strong acids, they all appear to have the same strength because they've all dissociated completely. And the only acid you end up with in the end is H3O+, which has the same strength from solution to solution. We call that the leveling effect. And so all these acids, they actually do have differing strengths here. But you can't tell the difference in water. You might have to use some sort of different solvent if you want to be able to figure out who's stronger and who's weaker and stuff like this. But in water, the leveling effect shows that they all appear to have the same strength because they all do dissociate completely giving H3O+. Now the strong bases are all going to be hydroxide salts, so metal hydroxides, and they're going to be the group one metal hydroxides, lithium hydroxide, sodium hydroxide, potassium hydroxide, rubidium hydroxide, and cesium hydroxide. And then most of the group two metal hydroxides, so magnesium hydroxide, calcium hydroxide, strontium hydroxide, and barium hydroxide. Now you might be like, well, Chad, why didn't, why didn't we include francium hydroxide on there? So, and that's a great question. And I'll, you know, I'll tell you what, you go test that out in the lab, you go play with some francium hydroxide. Let me know how that goes. And let me know if you find out if it's a strong base or not. So francium is super radioactive. You're never going to have a whole lot of it around. I think it has a half-life of like 22 minutes. So we don't really play with it, but if we did play with it, yes, it would end up on our strong base list, but you're never gonna play with it. So that's why we leave it off the list. So, and it turns out for the group 2 metal hydroxides, so it's a little bit tricky, because it turns out these aren't the most soluble things in the world. So it turns out as strong bases go, these actually, as you go down the periodic table, these do get stronger. And so like barium hydroxide is stronger than strontium hydroxide, which is stronger than calcium and stuff. That's not something you really probably need to know at this stage of the game. you want to just treat them all as strong. However, if you actually was working, you know, we're working with these in the lab, you find out that the group twos, they're not very soluble. Like maybe, you know, barium hydroxide, maybe somewhere around 0.1 molar. And, you know, calcium hydroxide, maybe somewhere around 0.01 molar. That's as much as you get to dissolve. But the amount that does dissolve will dissociate completely. So that's something to keep in mind. We'll revisit H2SO4 and these group two metal hydroxides in a little bit. The important part we're going to find out is that... H2SO4 is going to donate that first H 100%, but the second one is only weakly acidic and only partially going to dissociate. And we'll find out that the difference between strong and weak, strong acids dissociate effectively 100%, weak acids dissociate far less, partial dissociation, and usually less than 5%, and quite often like a lot less than 5%, like 0.01% is not uncommon and things of this sort. But it turns out it's not like a set number because it turns out it totally is concentration dependent. For a weak acid, the higher the concentration you have, the less it's going to dissociate. The lower the concentration you have, the more it's going to dissociate. I will bring that up again. It'll make a little more sense a little later on. We start doing pH calculations for weak acids. So, but just something to kind of file away in your head for now, and we'll talk about a little more later. Okay, so these are your strong acids, strong bases. You need to memorize these. If I give you any acid that's not one of these seven, you should automatically identify it as a weak acid. So if I say HF, well, the first thing you should realize, hopefully the pattern is clear, is that if something starts with an H, it's typically an acid. And in this case, HF is an acid. It starts with an H. but it's not on my strong acids list, so it must be a weak acid by default. And so now you should be able to identify any strong acid or any weak acid simply because you've memorized what the strong acids look like, and any other acid is therefore by default going to be weak. This is important too, because a lot of people think, oh, HF is super strong. No, it's not. Super deadly, it turns out, but it's because of the fluoride part of the molecule. You know, if you get a little HF on your skin, so... wash it off super fast, maybe just take a trip down to the ER just to be on the safe side, have them check you out. So HF will start going and dissolving through your skin and get into your bloodstream. It'll start leaching the calcium out of your blood. And calcium in your blood is kind of super important for like, you know, making sure you continue to have a heartbeat and things of the sort. So kind of important. So super deadly, but it's not because it's a strong acid. It is only a weak acid. And a lot of people, a lot of students know they've seen like Breaking Bad. They're like, oh, HF, that's nasty stuff. You can dissolve dead bodies with that. Yes, you can. But again, it's not because it's a strong acid. It's actually the fluoride part of it that's responsible for these things. So just want to make clear the air here. HI, HBR, HCL, all strong acids. HF is a weak acid. One other weak acid you should be familiar with is acetic acid. And recognizing him as an acid is tricky because he doesn't start with an H in the formula. It ends with an H. However, not every compound you see ending with an H is going to be an acid. It turns out the thing you want to find is the COOH, and sometimes that'll be written as CO2H. So a couple different ways to write acetic acid. And it doesn't have to be with a CH3 on the front, it could be a lot of different things. This is an example of what we call an organic acid, a carbon-based acid. So, in an organic acid, in this case what we call a carboxylic acid, always ends with this part, COOH, or CO2. So don't just think that any organic compound out there that ends with an H is going to be an acid. Nope. It's got to have this COOH at the end, but you should probably also be able to identify those as weak acids. That's one example of something that doesn't start with an H, but is still going to be a weak acid. Now, weak bases, on the other hand, they're a little more of the pain in the butt. So it's not sufficient for you to say, well, I know what all the strong bases look like, so then what do the weak bases look like? Well, that's going to be a little bit trickier. But because it's a little bit trickier, it's going to be a little harder to require you to understand what they look like. And so real common weak base is ammonia. And so ammonia is a weak base and you're going to want to file that away in your head. Definitely weak base and anything that kind of resembles ammonia in its structure is probably going to be a weak base as well. So if you notice ammonia. Looks like so. And it turns out we can have what are called organic amines related to ammonia structure, where I just start replacing these hydrogens with some carbon chains. So in fact, I don't like the way I've drawn that. We're gonna redraw that. It's actually bonded through the carbon, so I should write it that way. So, and this would have the formula CH3NH2 now. And you can kind of see the resemblance to the formula for ammonia. And these nitrogen containing compounds, whether it be pure ammonia or something like this guy, which is methyl amine, kind of an ammonia derivative, they are your classic weak bases. And they're not the only ones, but it's gonna be hard for us to give you a pattern to memorize, which actually is a favor for you, because having you recognize weak bases outside of ammonia or something looking very similar to ammonia, it's probably not something you're gonna be on the hook for. Cool. Those are your strong acids. In the subsequent lessons here, we're going to get into the pH scale and pH calculations, and we'll learn different ways for calculating the pH for strong acid bases versus weak acid bases. So being able to identify them is going to be super important going forward. Now if you found this lesson helpful and you want to support the channel, hit that thumbs up button. Best thing you can do to make sure YouTube shares this lesson with other students as well. And if you're looking for practice questions for acid bases or anything else in general chemistry, then check out my general chemistry master course. I'll be sure to leave a link in the description. It includes over 1200 practice questions, final exam rapid reviews, practice final exams. So if any of those interest you, check out my master course. Again, links in the description. A free trial is available. Happy studying.

I'm releasing several lessons a week throughout the school year, so if you want to be notified every time I post one, subscribe to the channel, click the bell notification. Okay, in this introduction to acids and bases, we're going to... Define acids and bases. Well, it turns out we're going to look at three different definitions, and we're going to use one of those definitions to talk about what are called conjugate acids and conjugate bases.

And then finally, we'll polish this lesson off by distinguishing between strong acids and bases and weak acids and bases. So let's start off with those three definitions. So the three definitions for acids and bases we've got are Arrhenius, Bronsted-Lowry, and Lewis.

And we're going to kind of go in historical order here, as well as this kind of order of progression where they get to be a more complete definition. And what I mean by this is that Not all acids and bases will actually satisfy Arrhenius'definition. He would not have considered all acids or bases that we now know of to be acids or bases. Now, Bronson Lowry's definition is a little broader and includes a greater number of acids and bases, but Lewis has the ultimate definition. Every acid or base that we currently, in modern times, identify as acids or bases, Lewis'definition will include all of them, but the other two, not quite so much.

Let's talk about this for a second. So we start with Arrhenius here. So Arrhenius first off, he only dealt with aqueous solutions.

That's it. And so everything had to be in water for Mr. Arrhenius here. And the way he looked at an acid is an acid is anything that causes the H3O plus concentration to go up in water. That's what he looked at. So let's take a look at what that ultimately means here.

So Most famous acid. I'm gonna ask you, what is the most famous acid? And you're not supposed to say LSD.

It is HCl, hydrochloric acid, found in lots of chemistry labs, so, but also found in your stomach. It's the same stuff that you put in your pool that might also be called muriatic acid. So it is the most famous acid, not LSD.

In this case, if you look at the reaction of what happens when you put HCl in water, so... HCl donates one of its H plus ions over to water to form H3O plus. So we call H3O plus hydronium.

And so it turns out that anything that when you place it in water causes you to form hydronium, that's what Arrhenius looked at as an acid here. And so in this case again, increases the concentration of hydronium. You might also consider this me, you know, something kind of like an H plus donor in water. But the key is it had to be in water for Arrhenius to be considering this.

Now if we look at bases on the other hand, So bases instead increase the concentration of hydroxide. Instead, the most famous base is NaOH, sodium hydroxide, and it simply dissociates to produce OH-hydroxide ions. Now it turns out when you've got an aqueous solution, so usually you have an equal concentration of H3O+, and OH-. We'll study this a little more in subsequent lessons in this chapter. So we've got equal amounts.

So, but it turns out if you add an acid, you're going to end up with more H3O plus than OH minus. And if you add a base, then you're going to end up with more OH minus than H3O plus. And there's this interplay between the two. So, but that's what Mr. Ranius considered here. Whereas Mr. Bronsted-Lowry came along and said, hey, we've got a little bit better definition.

So for one, this doesn't have to only be about water. We can use any solvent. We can use alcohols and things of this sort.

In fact, we don't even have to use a solvent. We can do acid-base reactions in the gas phase. So, and the way they looked at this then, is that an acid was simply an H plus donor.

Didn't have to be in water at all. It's just an H plus donor. And another word for H plus is a proton. So, proton donor, if you notice that hydrogen has one proton in its nucleus and then one electron.

And if you lose the electrons that you become H plus, then all you are is that proton in the nucleus. And so that's why H plus is often referred to not simply as the hydrogen ion, but also just simply as a proton. And so another way to say this is proton donor here for a Bronsted-Lowry acid.

That's the way they looked at it. On the base side, they kind of said, well, you know, rather than looking, you know, being about hydroxide, which would be, you know, kind of all about aqueous solutions, it's really just about this transfer of this H plus ion of this proton. And so, whereas the acid is the proton donor, the base is the proton acceptor.

And so we're really just looking at the transfer of protons. So if we go back to this, reaction with aqueous solutions. So Bronsted and Lowry would look at this and say, okay, between the two reactants, who gave away the H plus? Well, that would be the HCl.

So they define him as the acid. And who accepted the H+, who gained the H+, well, H2O turns into H3O+, gaining not just another hydrogen, but a plus charge at the same time, and that makes him the base. And so, again, this is happening in water, and so it's, you know, satisfying Arrhenius'definition of an acid-base reaction, stuff like that. So, however, it's also satisfying Bronsted-Lowry's.

However, again, Bronsted-Lowry says, you know, this, you know, The whole idea of acid bases doesn't have to involve water. And so in fact, we could mix HCl and ammonia in the gas phase. These could be gases, not aqueous solutions at all.

And you're still going to get a reaction going on here. HCl is still the acid. And in this case, ammonia would be the base. And so we take a look at this and who's the proton donor? Well, again, that's going to be the HCl.

He's giving away an H plus ion to become just plain old Cl minus. And who's accepting that or receiving it? NH3 turning into NH4 plus in this case, so he is the base. And so I just want you to see that all the acid-base reactions that take place in water that Arrhenius would have called acid bases, Bronsted and Lowry will also call acid and bases.

However, even the reactions that don't involve water, so Bronsted and Lowry are now gonna consider those acid-base reactions as well. And again, it's just the idea that we're just transferring a proton, follow that proton. So who gives away the H plus and who receives or accepts the H plus is how we'll identify acid base.

And it turns out this is the most useful definition, kind of working definition for identifying acid bases. So it's much easier to follow the H than what we're going to see with Mr. Lewis, his idea in a second. So when I'm having you identify acid bases throughout this chapter, most of the time what I'm expecting you to do is use Bronsted and Lowry's definition to do so. All right, so Mr. Lewis came along and said, hey, you know, your definition is really good, Mr. Bronson and Lowry, but I've got a few acids that you're not gonna classify as acids that I still will. And so if we take a look, he said, it's not about the protons, transfer protons, all about electrons.

And let's not put that in blue. And so rather than an acid being a proton donor, it's actually that an acid is an electron acceptor. And rather than a base being a proton acceptor, it is an electron donor.

And sometimes you'll hear electron pair acceptor and electron pair donor. And the idea is that we're making a new bond in an acid-base reaction. So, and as long as that new bond is being made to a hydrogen ion, an H plus ion, well then you'd consider that acid-base reaction not only a Lewis acid and base reaction, but a Bronsted-Lowry acid-base reaction, no big deal.

But if the new bond being made is to something other than hydrogen, then Bronsted-Lowry's definition won't apply, but Lewis's still will. So let's take a look at this here for a second. Let's take a look at this last one. Let's make some room. Now let's look at this instead from Lewis's perspective.

So Mr. Lewis here, it helps if we actually draw some Lewis structures. Those are the three that we're going to need here to take a look at. So, and you want to look and identify where you've made a new bond in the reaction, and that's right there in our case.

That's the new bond. And that new bond is between nitrogen and hydrogen. It's between this nitrogen right here and the hydrogen that was donated by HCl, again at least from Bronsted and Lowry's perspective. And so the question is, where did the electrons come from to make that bond?

Did they come from nitrogen or did they come from hydrogen? Well, they came from nitrogen. Nitrogen uses its lone pair of electrons.

It donates those electrons to be the ones that nitrogen and hydrogen will share. And so in that sense, he's the electron pair donor. And that, according to Mr. Lewis, would make him the base.

And HCL says, hey, I didn't contribute anything to make this bond. I just accepted your offer to make me a bond. So I'm the electron pair acceptor. Thank you very much. Cool.

That makes him the Lewis acid. And so if you classify anything as an acid or a base, Under any of the less complete definitions, they will still be classified as acids or bases under Mr. Lewis's definition, but sometimes it's a little harder to see. Notice I actually had to draw the Lewis structures to kind of figure out who's the Lewis acid and who's the Lewis base.

Now there's also a class of acids that are just going to act as Lewis acids, but they're not going to be a part of any other definition. Arrhenius wouldn't have considered them necessarily acids or Bronsted-Lowry wouldn't consider acids, but Lewis has a small class of acids that he would consider acids. We want to take a look at those as well. In fact, we'll just do that right over here.

So if we take a look, boron compounds are really commonly gonna fall into this category. I'm gonna leave off a lot of the lone pairs here that are on the fluorines. We don't need those for this reaction. And so boron, you might recall, is famous for not having a filled octet.

And not having a filled octet, boron actually has an empty orbital. And an empty orbital makes you a good candidate for being an electron acceptor. It turns out, a place where you can accept those electrons.

And so boron-containing compounds are often going to be able to act as Lewis acids, and only going to be acting as a Lewis acid by any of the definitions. Won't satisfy either the Uranus or Bronsted-Lowry definition. definition, but also metal ions, it turns out.

Metal ions, being cations, have lots of empty orbitals in their valence shell quite commonly, and so metal ions will often act as electronic scepters, Lewis acids as well, even though, again, Arrhenius and Bronsted-Lowry would not have necessarily considered them acids. So, just want to give you an idea of, you know, if I ask you to identify an acid that only satisfies Lewis's definition only, so... probably going to be a boron containing compound or a metal ion like aluminum ions are a good example. And so if we look at what actually happens here in this reaction, we end up with a positive formal charge of nitrogen and negative formal charge on boron. And the new bond we made is the bond between nitrogen and boron.

And the question is, well, where did those electrons come from? From nitrogen or from boron? Well, they're again coming from the lone pair on the nitrogen. Boron didn't have a lone pair to donate here. And so the hallmark here again of our Lewis acid, it's going to have empty orbitals.

And that's usually again going to be, if it's going to be uniquely a Lewis acid, it's going to be a boron containing compound or a metal ion. But we also saw that, you know, HCl satisfied, you know, the qualifications for being a Lewis acid as well. If you're an Arrhenius acid or Bronsted-Lowry acid, then you are also a Lewis acid. Again, Lewis had the ultimate definition. Now, the hallmark of being the Lewis base here, and you should recognize this, is to act as a Lewis base, you need a lone pair of electrons.

That is the one qualification you must have. Cool. And as a Lewis base, if you use that lone pair of electrons to make a new bond to a hydrogen, again, then not only are you a Lewis base, you're going to be a Bronsted-Lowry base as well.

That's kind of the distinguishing factor. But if you use that lone pair of electrons to make a bond to any other atom besides hydrogen like boron, then you're going to be a Lewis base, but you're not going to be a Bronsted-Lowry base in that case. So those are your three definitions.

You should definitely know verbatim what these definitions are, but you probably want to know how to apply them as well. And again, as far as working definition for identifying acids and bases, We're going to use Bronsted-Lowry's definition more commonly than any other. However, you should keep in mind that Lewis has the ultimate definition. If you are an acid, there's a good chance you're probably an Arrhenius acid.

There's a really good chance that you're probably a Bronsted-Lowry acid, but there is a 100% chance that you're a Lewis acid. He has the most complete definition. All right, so I left the Bronsted-Lowry definitions for acid bases on the board here because...

That's going to be crucial for us identifying what are referred to as conjugate acid conjugate base pairs. And so, it turns out this is going to have some relevance when we start talking about trends in acidity and trends in basicity and stuff like this. And it might just really be a question on the test that you have to identify an acid's conjugate base or a base's conjugate acid. And so it turns out you've often got these two species in equilibrium, a conjugate acid conjugate base pair. in equilibrium together in most solutions and stuff like this.

You might have more of one than the other depending on conditions and stuff like this, but they're usually in equilibrium together, and their relationship often determines how strong the acid is or how strong the base is, things of this sort. So super important concept to identify, and it is Bronsted-Lowry's definition that's going to allow us to identify these. So if you take a look at HCl here, I want to fill in this table here, and we want to know what is HCl's conjugate base. So here's the conjugate acid.

The question you get on the test might just be, What is the conjugate base of HCl? Well what you want to do here is you just want to have HCl act as an acid. Well what do acids do according to Mr. Bronson-Lowry? Well they're H plus donors.

So he's going to donate an H plus and whatever you have left when he's done, that is your conjugate base. Well after he gets rid of the H plus, he has one fewer H's. So now we have a chlorine atom with no H's, only had one to start with, and one fewer plus charge.

Well, he was a neutral species before, and so if you lose a positive charge, you're now going to have a negative charge. And so Cl-would be properly identified as the conjugate base of HCl. Cool. And the big thing is that after HCl, which again, most famous strong acid, after he acts as an acid, what he turns into is his conjugate base. Cool, let's go the other way here.

So we've got ammonia next on the chart listed as a conjugate base. And the question you might see on a test is which of the following is the conjugate acid of ammonia. So we'll start with ammonia here.

And in this case, we just need him. He's been identified as the base here because we're looking for his conjugate acid. He must be the base.

And so we just need him to act as a base. Well, what do bases do? Well, they're H plus acceptors.

So we just need him to accept. an H+. And so if he accepts an H+, meaning he gains an H+, so instead of having three of those H's, he's now going to have four, but he also gains a positive charge as well. So he was a neutral molecule before, he's now going to have an overall plus one charge. And that NH4 plus would be the conjugate acid here we'd fill in on the chart.

And so the question again, if the question was, what is the conjugate acid of NH3? you'd say NH4+. Okay, so now we've got to get a little bit tricky here. And the way we're going to get tricky is by introducing what we're going to call amphiprotic.

So you might think of ampha like an amphibian as an animal that can live on land or water. Well, amphiprotic is a species that can act as a proton donor or a proton acceptor. it can act as a Bronsted-Lowry acid or a Bronsted-Lowry base.

That's what amphiprotic means. And we're going to choose such compounds because if I asked you to identify a conjugate acid or a conjugate base, well, amphiprotic species have both a conjugate acid and a conjugate base. And so if you don't realize which one you're being asked or get it backwards in your mind, it's a great question to kind of slip you up on the test. And so in this case, we're going to start with bicarbonate here, HCO3-. And so it turns out bicarbonate can act as an acid and donate another H plus ion.

or it can act as a base and accept an H plus ion. And so you gotta really know your definitions here so you can figure this out. So if we start with bicarbonate here and we say, what is the conjugate base of bicarbonate?

What goes in this box? Well, if I'm asking for the conjugate base, then it can be implied that the bicarbonate itself must be the conjugate acid form. And what do acids do?

Acids donate H plus, they lose that H plus. And so if we lose an H plus from the formula here, then we'll have no Hs left. And again, losing a plus charge, well, we're already negative one, so we're going to be one less positive, i.e. one more negative, and so we'll have a negative two charge. And so we get the carbonate ion.

That is the conjugate base of the bicarbonate ion. But you could ask this backwards. Again, bicarbonate is amphiprotic. It can act as an acid or base.

And so if instead, we want to know what goes in this box. What is the conjugate acid of the bicarbonate ion of HCO3-? Well, again, if I want the conjugate acid, then I got to know that HCO3-is acting as the base. He's the conjugate base this time.

What do bases do? Well, they're H plus acceptors. They gain H plus.

And so in this case, if HCO3-gains an H, he's now going to have two of them. But he also gains a plus charge again. And so if he started out as minus one, he's now going to be neutral. And H2CO3 there, carbonic acid, is the conjugate acid of HCO3-bicarbonate.

Now, if I want to get super tricky on this, I will do the same thing with hydroxide. We normally don't think of hydroxide as being amphiprotic. Hydroxide is pretty much like your classic base. In an aqueous solution, hydroxide is kind of how you recognize you have base in the solution. And so whenever you hear hydroxide, it's going to be really kind of this association in your mind that it's a base.

And so if I said, hey, what's the conjugate of hydroxide? You're not even going to think like, oh, you didn't. You missed a word there, Chad. You just said conjugate.

You didn't say conjugate acid or base. Which one did you mean? You're just going to think, oh, he must have just meant conjugate acid. Well, and that could be the question.

What is the conjugate acid of hydroxide? Well, if I'm looking for the conjugate acid, then hydroxide must be acting as the conjugate base. And what do bases do? They're H plus ion acceptors.

And if OH minus accepts an H, you're now going to have two Hs with that oxygen. And it's going to gain a plus charge to now be neutral. And that's water.

And so water is the conjugate acid of hydroxide. But if I was trying to be tricky because, you know, maybe my team just lost the Super Bowl and I'm angry at all my students, or at least I'm going to take it out on them. So then I might ask, what's the conjugate base of hydroxide? Well, we don't normally see hydroxide acting as an acid, but if I'm asking for his conjugate base, then he must be the conjugate acid. This is super uncommon.

It turns out you can get a hydroxide to act as an acid if you react it with an even stronger base, which is rare and not something you're likely to encounter in this class, but in organic chemistry you might. So, but hydroxide, if he's the acid and I'm looking for his conjugate base, well then have him act as the acid, have him donate another H plus ion. And if he gets rid of an H, then he won't have any left, that oxygen, and he'll also lose a plus charge.

So instead of being minus one, he's now going to be minus two and you get the oxide ion. And so the conjugate base of hydroxide is the oxide anion here, O2 minus. Super tricky.

I've seen it show up a time or two on an exam. Only reason I'm covering it. Cool.

So hopefully now you can identify conjugate acids and conjugate-based pairs. And if I give you one, you can ask for the other and things of this sort. You might also want to be able to do one other thing. Alright, so I'm going to try and give you an example that's not so obvious.

And you might have noticed I've been using a lot of red and blue. And red for acids and blue for bases. So I'm not doing that here. So, and if you don't know what this is from, it's from litmus paper.

Litmus paper turns red in the presence of acids, and it turns blue in the presence of bases. It's a nice little quick test to see if you've got an acid or a base. Well, here I'm giving this to you all in black, so you can't figure it out. So, and I'm giving you an acid-base reaction that is reversible, and most acid-base reactions, it turns out, are reversible.

And so I want you to identify the Bronsted-Lowry acids and bases in this reaction. And you can look at it going in the forward direction, but it is reversible, and you can look at it going in the reverse direction. And so what you're going to find is in a typical acid-base reaction, according to Mr. Bronsted and Lowry here, you're going to have an acid base on the reactant side going in the forward direction, but you'll have an acid base in the product side going in the reverse direction, and you just got to figure out which based on your definitions.

And so if we look at this in the forward direction, who is gaining the H plus ion? Well in this case that's NH3. He's going from NH3 to NH4, and we can see that CH3COOH is the one losing the H, and it's that last H plus ion that he's losing.

in the process. And so the one gaining it, the H plus acceptor, is our base. And the one losing it is our acid.

But notice, identifying the ones on the other side, well, in this case, again, if ammonia here is the base, well, after he acts as a base, what do you turn into? Well, after you act as your base, you turn into your conjugate acid. And so on this side, if I was doing the reverse direction, NH4 plus would be the one donating the H plus.

And I see my conjugate base conjugate acid pair right there. So, and same thing the other way, if I was doing the reverse reaction, it'd be CH3COO minus, that would be gaining the H. And he's the conjugate base of this acid right here.

Cool, and so in a typical Bronsted-Lowry acid-base reaction, you should be able to identify the acid and base both on the reactant and the product side as well. All right, so the last topic in this lesson, we've got to talk about the difference between strong acids and bases and weak acids and bases. And the difference in strength, it turns out, is how much they dissociate.

So if you're a proton donor, a Bronsted-Lowry acid, it's how much you donate that proton. That's what we're going to refer to as being as dissociation or ionization. All right, so if we take a look, we've got seven strong acids on our list. I will let you know that most of you are either going to have six or seven.

So, and that last one, HClO3, it shows up on some lists, and it doesn't show up on others. It's kind of like a, right at the cusp of being strong, and some people put it on their list, and some people don't. Now, there are not only seven strong acids in the universe, but there's not too many more, it turns out.

But these are the seven you're likely to see, again, if you're not, in some cases, going to be given only six. And it turns out, if you see an acid that's any other than these, you should identify it as a weak acid, which means it will not dissociate 100%. So...

strong acids associate 100% in water. So if we go back to our most famous acid again, HCl, we saw its reaction with water earlier to form hydronium and its conjugate base, and this effectively goes 100% to completion. Now truth be told it's not 100%, but it's like 99.99%, so it's effectively 100% to completion.

And the truth is that would happen for any of these acids. So if you put any of these acids in water, they disappear completely because they dissociate completely. And the only acid that will actually be present in that solution will be H3O+.

And so it turns out when you have aqueous solutions of any of the strong acids, they all appear to have the same strength because they've all dissociated completely. And the only acid you end up with in the end is H3O+, which has the same strength from solution to solution. We call that the leveling effect. And so all these acids, they actually do have differing strengths here.

But you can't tell the difference in water. You might have to use some sort of different solvent if you want to be able to figure out who's stronger and who's weaker and stuff like this. But in water, the leveling effect shows that they all appear to have the same strength because they all do dissociate completely giving H3O+. Now the strong bases are all going to be hydroxide salts, so metal hydroxides, and they're going to be the group one metal hydroxides, lithium hydroxide, sodium hydroxide, potassium hydroxide, rubidium hydroxide, and cesium hydroxide. And then most of the group two metal hydroxides, so magnesium hydroxide, calcium hydroxide, strontium hydroxide, and barium hydroxide.

Now you might be like, well, Chad, why didn't, why didn't we include francium hydroxide on there? So, and that's a great question. And I'll, you know, I'll tell you what, you go test that out in the lab, you go play with some francium hydroxide. Let me know how that goes.

And let me know if you find out if it's a strong base or not. So francium is super radioactive. You're never going to have a whole lot of it around. I think it has a half-life of like 22 minutes. So we don't really play with it, but if we did play with it, yes, it would end up on our strong base list, but you're never gonna play with it.

So that's why we leave it off the list. So, and it turns out for the group 2 metal hydroxides, so it's a little bit tricky, because it turns out these aren't the most soluble things in the world. So it turns out as strong bases go, these actually, as you go down the periodic table, these do get stronger. And so like barium hydroxide is stronger than strontium hydroxide, which is stronger than calcium and stuff. That's not something you really probably need to know at this stage of the game.

you want to just treat them all as strong. However, if you actually was working, you know, we're working with these in the lab, you find out that the group twos, they're not very soluble. Like maybe, you know, barium hydroxide, maybe somewhere around 0.1 molar.

And, you know, calcium hydroxide, maybe somewhere around 0.01 molar. That's as much as you get to dissolve. But the amount that does dissolve will dissociate completely.

So that's something to keep in mind. We'll revisit H2SO4 and these group two metal hydroxides in a little bit. The important part we're going to find out is that...

H2SO4 is going to donate that first H 100%, but the second one is only weakly acidic and only partially going to dissociate. And we'll find out that the difference between strong and weak, strong acids dissociate effectively 100%, weak acids dissociate far less, partial dissociation, and usually less than 5%, and quite often like a lot less than 5%, like 0.01% is not uncommon and things of this sort. But it turns out it's not like a set number because it turns out it totally is concentration dependent. For a weak acid, the higher the concentration you have, the less it's going to dissociate. The lower the concentration you have, the more it's going to dissociate.

I will bring that up again. It'll make a little more sense a little later on. We start doing pH calculations for weak acids.

So, but just something to kind of file away in your head for now, and we'll talk about a little more later. Okay, so these are your strong acids, strong bases. You need to memorize these. If I give you any acid that's not one of these seven, you should automatically identify it as a weak acid.

So if I say HF, well, the first thing you should realize, hopefully the pattern is clear, is that if something starts with an H, it's typically an acid. And in this case, HF is an acid. It starts with an H. but it's not on my strong acids list, so it must be a weak acid by default. And so now you should be able to identify any strong acid or any weak acid simply because you've memorized what the strong acids look like, and any other acid is therefore by default going to be weak.

This is important too, because a lot of people think, oh, HF is super strong. No, it's not. Super deadly, it turns out, but it's because of the fluoride part of the molecule.

You know, if you get a little HF on your skin, so... wash it off super fast, maybe just take a trip down to the ER just to be on the safe side, have them check you out. So HF will start going and dissolving through your skin and get into your bloodstream.

It'll start leaching the calcium out of your blood. And calcium in your blood is kind of super important for like, you know, making sure you continue to have a heartbeat and things of the sort. So kind of important.

So super deadly, but it's not because it's a strong acid. It is only a weak acid. And a lot of people, a lot of students know they've seen like Breaking Bad.

They're like, oh, HF, that's nasty stuff. You can dissolve dead bodies with that. Yes, you can. But again, it's not because it's a strong acid.

It's actually the fluoride part of it that's responsible for these things. So just want to make clear the air here. HI, HBR, HCL, all strong acids.

HF is a weak acid. One other weak acid you should be familiar with is acetic acid. And recognizing him as an acid is tricky because he doesn't start with an H in the formula. It ends with an H.

However, not every compound you see ending with an H is going to be an acid. It turns out the thing you want to find is the COOH, and sometimes that'll be written as CO2H. So a couple different ways to write acetic acid.

And it doesn't have to be with a CH3 on the front, it could be a lot of different things. This is an example of what we call an organic acid, a carbon-based acid. So, in an organic acid, in this case what we call a carboxylic acid, always ends with this part, COOH, or CO2. So don't just think that any organic compound out there that ends with an H is going to be an acid. Nope.

It's got to have this COOH at the end, but you should probably also be able to identify those as weak acids. That's one example of something that doesn't start with an H, but is still going to be a weak acid. Now, weak bases, on the other hand, they're a little more of the pain in the butt.

So it's not sufficient for you to say, well, I know what all the strong bases look like, so then what do the weak bases look like? Well, that's going to be a little bit trickier. But because it's a little bit trickier, it's going to be a little harder to require you to understand what they look like. And so real common weak base is ammonia.

And so ammonia is a weak base and you're going to want to file that away in your head. Definitely weak base and anything that kind of resembles ammonia in its structure is probably going to be a weak base as well. So if you notice ammonia. Looks like so. And it turns out we can have what are called organic amines related to ammonia structure, where I just start replacing these hydrogens with some carbon chains.

So in fact, I don't like the way I've drawn that. We're gonna redraw that. It's actually bonded through the carbon, so I should write it that way. So, and this would have the formula CH3NH2 now.

And you can kind of see the resemblance to the formula for ammonia. And these nitrogen containing compounds, whether it be pure ammonia or something like this guy, which is methyl amine, kind of an ammonia derivative, they are your classic weak bases. And they're not the only ones, but it's gonna be hard for us to give you a pattern to memorize, which actually is a favor for you, because having you recognize weak bases outside of ammonia or something looking very similar to ammonia, it's probably not something you're gonna be on the hook for.

Cool. Those are your strong acids. In the subsequent lessons here, we're going to get into the pH scale and pH calculations, and we'll learn different ways for calculating the pH for strong acid bases versus weak acid bases. So being able to identify them is going to be super important going forward.

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Happy studying.

Transcript for:Understanding Acids and Bases in Chemistry

Transcript for:
Understanding Acids and Bases in Chemistry