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0:01
- An introduction to acids and bases gonna 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 where you can find those courses. Now, this lesson's part of my new general chemistry playlist. I'm releasing several lessons a week throughout the school year.
0:22
So if you wanna 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 gonna define acids and bases, and it turns out we're gonna look at three different definitions, and we're gonna 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.
0:45
So let's start off with those three definitions. So the three definitions for acids and bases, we've got our Arrhenius, Bronsted-Lowry, and Lewis. And we're gonna 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's definition. He would not have considered all acids or bases that we now know of to be acids or bases.
1:08
Now, Bronsted and Lowry's definition's 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's definition will include all of them. But the other two, not quite so much. Let's talk about this for a second. So if we start with Arrhenius here, so Arrhenius first off, he only dealt with aqueous solutions, that's it.
1:37
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
1:48
H3O+ 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, 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
2:16
of what happens when you put HCl in water, So HCl donates one of its H+ ions over to water to form H3O+.
2:26
So we call HCl+ 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 making something kind of like an H+ 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
2:56
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+
3:14
and OH-. We'll study this a little more in subsequent lessons in this chapter. So if you've got equal amounts, but it turns out if you add an acid, you're gonna end up with more H3O+ than OH-. And if you add a base, then you're gonna end up with more OH- than H3O+, and there's this interplay between the two. But that's what Mr. Arrhenius considered here. Whereas Mr. Bronsted-Lowry came along and said, hey, we've got a little bit better definition.
3:36
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. And the way they looked at this then is that an acid was simply
3:54
an H+ donor, didn't have to be in water at all. It's just an H+ donor, and another word for H+ 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 electron, that you become H+, then all you are is that proton in the nucleus. And so that's why H+ is often referred to, not simply as the hydrogen ion, but also just as something that has a proton. And so another way to say this is proton donor here
4:20
for a Bronsted-Lowry acid. That's the way they looked at it. On the base side, they kind of said, well, rather than being about hydroxide, which would be kind of all about aqueous solutions, it's really just about this transfer of this H+ ion of this proton. And so, whereas the acid is the proton donor, the base is the proton acceptor.
4:42
And so we're really just looking at the transfer protons. So if we go back to this reaction with aqueous solutions, so Bronsted and Lowry look at this and say, okay, between the two reactants, who gave away the H+? Well, that would be the HCl.
4:57
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
5:10
him the base. And so, again, this is happening in water. And so it's satisfying Arrhenius's definition of an acid-base reaction and stuff like that. So, however, it's also satisfying Bronsted-Lowry's. However, again, Bronsted-Lowry says this whole idea of acid bases doesn't have to involve water. And so, in fact, we could mix HCl,
5:32
and ammonia in the gas phase. This could be gases, not aqueous solutions at all. And you're still gonna get a reaction going on here. HCl is still the acid, and in this case, ammonia would be the base.
5:46
And so we take a look at this, and who's the proton donor? Well, that again, that's gonna be the HCl. He's giving away an H+ ion to become just plain old Cl-. And who's accepting that or receiving it? NH3, turning into NH4+ in this case, so he is the base.
6:03
So I just want you to see that all the acid-base reactions that take place in water that Arrhenius would have called acid and bases, Bronsted and Lowry will also call acid and bases. However, even the reactions that don't involve water, so Bronsted-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+ and who receives or accepts the H+ is how we'll identify acid-base.
6:27
And it turns out this is the most useful definition, kind of working definition, for identifying acids and bases. So it's much easier to follow the H then what we're gonna see with Mr. Lewis, his idea, in a second. So when I'm having you identify acids and 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.
6:50
All right, so Mr. Lewis came along and said, hey, your definition's really good, Mr. Bronsted 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, it's all about electrons, and let's not put that in blue.
7:10
And so rather than an acid being a proton donor, it's actually that an acid is an electron acceptor.
7:19
And rather than an 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. And as long as that new bond is being made to a hydrogen ion, an H+ 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
7:46
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.
7:59
Now, let's look at this instead from Lewis's perspective. So Mr. Lewis here, and it helps if we actually draw some Lewis structures.
8:23
Those are the three that we're gonna need here to take a look at.
8:27
And you wanna 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.
8:52
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
9:20
under any of the less complete definitions, they will still be classified as acids or bases under Mr. Lewis's definition, but sometimes a little hard 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?
9:34
Now there's also a class of acids that are just going to act as Lewis acids, but they're not gonna be a part of any other definition. Arrhenius wouldn't have considered them necessarily acids, or Bronsted-Lowry wouldn't consider them acids, but Lewis has a small class of acids that he would consider acids. We wanna take a look at those as well. In fact, we'll just do that right over here.
9:57
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.
10:13
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 gonna be able to act as Lewis acids and only gonna be acting as a Lewis acid by any of the definitions. Won't satisfy either the Arrhenius or Bronsted-Lowry definition,
10:41
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 electron acceptors, Lewis acids as well, even though again, Arrhenius and Bronsted-Lowry would not have necessarily considered them acids. So just wanna give you an idea... If I ask you to identify an acid that only satisfies Lewis's definition only. So probably gonna be a boron-containing compound
11:07
or a metal ion, like aluminum ions are a good example. So if we look at what actually happens here in this reaction,
11:25
we end up with a positive formal charge on 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 gonna have empty orbitals. And that's usually, again, going to be,
11:48
if it's gonna be uniquely a Lewis acid, it's gonna be a boron-containing compound or a metal ion, but we also saw that HCl satisfied 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,
12:09
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.
12:20
Cool, and as a Lewis base, if you use that lone pair of electrons to make a new bond to a hydrogen, then not only are you a Lewis base, you're gonna 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 gonna be a Lewis base, but you're not gonna be a Bronsted-Lowry base in that case. So those are your three definitions. You should definitely know verbatim
12:47
what these definitions are. But you probably wanna know how to apply them as well. And again, as far as a working definition for identifying acids and bases, we're gonna 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,
13:11
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 and bases on the board here, because that's gonna be crucial for us identifying what are referred to as conjugate acid, conjugate base pairs. And so it turns out, this is gonna 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
13:36
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.
13:58
So super important concept to identify. And it is Bronsted-Lowry's definition that's gonna allow us to identify these. So if you take a look at HCl here, I wanna fill in this table here and we wanna know what is HCl's conjugate base? So he's the conjugate acid. The question you've got on the test might just be, what is the conjugate base of HCl? Well, what you wanna do here is you just wanna have HCl act as an acid. Well, what do acids do, according to Mr. Bronsted-Lowry?
14:24
Well, there they're H+ donors. So he's gonna donate an H+ and whatever you have left when he's done, that is your conjugate base. Well, after he gets rid of the H+ 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 gonna have a negative charge. And so Cl- would be properly identified as the conjugate base of HCl.
14:56
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.
15:20
And in this case, we just need him... He's been identified as the base here, 'cause 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+ acceptors. So we just need him to accept an H+. And so if he accepts an H+, meaning he gains an H+, well, instead of having three of those H's, he's now gonna have four, but he also gains a positive charge as well. So he was a neutral molecule before.
15:47
He's now gonna have an overall plus one charge. And that NH4+ 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+.
16:05
Okay, so now we've gotta get a little bit tricky here and the way we're gonna get tricky
16:17
is by introducing what we're gonna call amphiprotic. So you might think of amphi, 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. We're gonna choose such compounds because, if I asked you to identify conjugate acid or conjugate base, well, amphiprotic species have both
16:44
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 gonna start with bicarbonate here, HCl3-. And so it turns out bicarbonate can act as an acid and donate another H+ ion, or it can act as a base and accept an H+ ion. And so you gotta really know your definitions here so you can figure this out.
17:07
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+. They lose that H+. And so if we lose an H+ from the formula here, then we'll have no H's left, and again, losing a plus charge. Well, we're already negative one, so we're gonna be one less positive,
17:35
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 wanna 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 gotta know that HCl3- is acting as the base. He's the conjugate base this time.
18:05
Well, what do bases do? Well, they're H+ acceptors, they gain H+. And so in this case, if HCl3- gains an H, he's now gonna 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 HCl3- bicarbonate. Now, if I wanna get super tricky on this, I will do the same thing with hydroxide. We normally don't think of hydroxide as being amphipathic.
18:36
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 gonna 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 gonna think like, oh, 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 gonna think, oh,
19:02
he must've 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+ ion acceptors. And if OH- accepts an H, you're now gonna have two H's with that oxygen, and it's gonna 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,
19:29
because maybe my team just lost the Superbowl and I'm angry at all my students, so I'm gonna 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 its conjugate base, then he must be the conjugate acid. This is super uncommon. It turns out you can get 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
19:56
in this class, but in organic chemistry you might. But hydroxide, if he's the acid. I'm looking for his conjugate base. Well, then have him act as the acid. have him donate another H+ 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 gonna be minus two and you get the oxide ion. And so the conjugate base of hydroxide is the oxide anion here, O2-. Super tricky.
20:23
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 base pairs. And if I give you one, you can ask for the other and things of this sort. You might also wanna be able to do one other thing.
21:01
All right, so I'm gonna try and give 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. 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 got an acid or a base. Well here, I'm giving this to you all in black
21:21
so you can't figure it out. And I'm giving you an acid-base reaction that is reversible. And most acid-base reactions, 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 gonna find is in a typical acid-base reaction, according to Mr. Bronsted and Lowry here,
21:43
you're gonna have an acid and a base on the reactant side going in the forward direction, but you'll have an acid and 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+ 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+ ion that he's losing in the process.
22:09
And so the one gaining it, the H+ acceptor, is our base,
22:15
and the one losing it is our acid.
22:20
But notice that identifying the ones on the other side, well, in this case again, if ammonia here's 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+ would be the one donating the H+. And I see my conjugate base, conjugate acid pair right there. So same thing the other way, if I was doing the reverse reaction, it'd CH3COO- that would be gaining the H,
22:49
and he's the conjugate base of this acid right here.
22:57
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 gotta 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 gonna refer to
23:22
as being 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 are either gonna have six or seven. And that last one, HClO3, it shows up on some lists, and it doesn't show up on others. It's kind of like right at the cusp of being strong, and some people put 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,
23:47
but these are the seven you're likely to see. Again, if you're not, in some cases, gonna 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 dissociate 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.
24:19
Now, truth be told, it's not 100%, but it's like 99.999%, so it's effectively 100% to completion. And the truth is, that'd 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
24:45
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 wanna be able to figure out who's stronger and who's weaker and stuff like this. But in water, the leveling effect shows
25:07
that they all appear to have the same strength because they all do dissociate completely, giving H3O+. Now the strong bases are all gonna be hydroxide salts, so metal hydroxides, and they're gonna be the Group 1 metal hydroxides lithium hydroxide, sodium hydroxide, potassium hydroxide, rubidium hydroxide, and cesium hydroxide. And then most of the Group 2 metal hydroxides. So magnesium hydroxide, calcium hydroxide, strontium hydroxide, and barium hydroxide. Now you might be like, well, Chad,
25:32
why didn't we include francium hydroxide on there? And that's a great question. 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 gonna 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,
25:57
it would end up on our strong base list, but you're never gonna play with it. So that's why we'll leave it off the list. And it turns out for the Group 2 metal hydroxides, so this is a little bit tricky 'cause 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,
26:19
but that's not something you really probably need to know at this stage of the game. You wanna just treat them all as strong. However, if you actually were working with these in the lab, you'd find out that the Group 2s, they're not very soluble, like maybe barium hydroxide may be somewhere around 0.1 molar, and calcium hydroxide may be somewhere around 0.01 molar. That's as much as you get to dissolve. But the amount that does dissolve will dissociate completely.
26:42
So that's something to keep in mind. We'll revisit H2SO4, and these Group 2 metal hydroxides in a little bit, 'cause the important part we're gonna find out is that H2SO4 is gonna donate that first H 100%. But the second one is only weakly acidic, and only partially gonna dissociate. And we'll find out that the difference between strong and weak, strong acids associate effectively 100%, weak acids dissociate far less, partial dissociation, and usually less than 5%.
27:12
And quite often, like a lot less than 5%. Like 0.01% is not uncommon and things of this sort.
27:18
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. But just something to kind of file away in your head for now. And we'll talk about it a little more later.
27:43
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,
27:56
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 gonna be weak.
28:21
This is important, too, 'cause a lot of people think, oh, HF is super strong. No, it's not. Super deadly, it turns out, but it's 'cause 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 leeching the calcium out of your blood,
28:44
and calcium in your blood is kind of super important for like making sure you continue to have a heartbeat and things of this 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 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. (laughs) But again, it's not because it's a strong acid. It's actually the fluoride part of it
29:06
that's responsible for these things.
29:09
So just wanna 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 gonna be an acid. It turns out the thing you wanna find is the COOH, and sometimes that'll be written as CO2H.
29:42
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 CO2H.
29:59
So don't just think that any organic compound out there that ends with an H is gonna 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. Just one example of something that doesn't start with an H, but it's 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.
30:23
So then what are the weak bases look like? Well, that's gonna be a little bit trickier. Because it's a little bit trickier, it's gonna be a little harder to require you to understand what they look like. And so real common weak base
30:38
is ammonia. And so ammonia is a weak base, and you're gonna wanna file that away in your head. Definitely a weak base, and anything that kind of resembles ammonia in its structure is probably gonna be a weak base as well. So if you notice ammonia
30:58
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. And this would have the formula CH3NH2 now,
31:21
and you can kind of see the resemblance to the formula for ammonia.
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And these nitrogen-containing compounds, whether it be pure ammonia or something like this guy, which is methylamine, 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.
31:49
Cool, those are your strong acids. In the subsequent lessons here, we're gonna get into the pH scale and pH calculations, and we'll learn different ways for calculating the pH for strong acid and bases versus weak acids and bases, so being able to identify them is gonna be super important going forward.
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Now, if you found this lesson helpful, and you wanna 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 acids and 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 1,200 practice questions, final exam rapid reviews, practice final exams.
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So if any of those interest you, check out my master course. Again, links in the description. A free trial is available. Happy studying.
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