Carboxylic acid introduction
Carboxylic acid introduction. Created by Sal Khan.
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- How come oxygen gives an electron isn't it the one of the elements with the highest electronegative so why does it give away an electron(6 votes)
- It isn' t that the O atom wants to give away its electrons, but that it can (reluctantly) give them away. Actually, it doesn't give them away, it just gives the C atom a little more control over them. We have a tug of war between the carbonyl O and the negatively charged O, both of whom want electrons, but the C=O wants them more. So the molecule reaches a compromise in which each O atom gets half of the negative charge (the resonance hybrid). This resonance hybrid is more stable than the structure in which one O atom has a full negative charge and the other has none.
Remember that electrons repel each other. In the hybrid, half the electron density is on one O atom, and half the electron density is on the other O atom. This puts the electrons as far away from each other as possible, minimizes the repulsions, and lowers the energy of the carboxylate ion.(10 votes)
- At08:00, what are the two single- and one double-bond connected to? Each other? Or a carbon atom? Please answer quickly!(5 votes)
- It's implied that it is a carbon atom. When he talks about dot structures, Sal explains that, to make them, you ignore all hydrogen atoms (or rather, imply they are there) and draw the carbon bonds but don't write "C" every time. This makes it a lot quicker to jot down complex carbon molecules, because you don't have to write every "H" and "C." I can't find his video on this.(1 vote)
- i am not sure if this question is refered to the vedio but i am getting too much confused about this
RCOOH + NaHCO3 ----> RCOONa + CO2 + H2O
how the organic acid replace the Na from carbonic acid and it said that organic acids are weaker than inorganic acids so how organic acid replace the salt of inorganic acid?(4 votes)
- What does the R stand for? Is it where the chain he was referencing at0:20is supposed to go?(2 votes)
- The R is the general symbol for an alkyl group (methyl, ethyl, etc.)
No, he is talking about the OH group on the other side.(3 votes)
- Why does Acetic Acid is soluble in water?(2 votes)
- It's a bit of a balancing game.
The -CH3 bit is non polar, but the -COOH bit is quite polar and that is why it dissolves in water.
As the length of the alkyl chain gets bigger though, these carboxylic acids become less and less water soluble, eg. you generally wouldn't really say octanoic acid is water soluble, the non polar alkyl chain dominates now.(2 votes)
- Hi ! Are carboxylic acids more acidic than phenols as well.......and if they are, aren't they like more acidic than water?(2 votes)
- phin oxide have almost the same reason as carboxylates for stablization... then why carboxlylate is more stable the phin oxide ion(2 votes)
- Why does a more stable acid, here a resonant ion, mean greater "acidity"? Shouldn't structures which are more likely to react be more acidic? Thanks.(2 votes)
- Although phenoxide has more number of resonating structures than carboxylate ion, carboxylic acid is a stronger acid than phenol. why?(1 vote)
- Resonance structures involving the ring disrupt the cyclic 6π system, so these structures are minor contributors.
The resonance contributors to the carboxylate ion are equally important.
Carboxylate ion has more resonance stabilization than phenoxide ion, so the position of equilibrium lies further to the right.
Hence, a carboxylic acid is a stronger acid than phenol.(1 vote)
- water is more acidic than carboxylic acids?(1 vote)
Let's see if we can learn thing or two about carboxylic acids. They have the general form of the carbonyl group, just like we've seen in aldehydes and ketones, and they will be part of a longer carbon chain. But instead of having a hydrogen here, as the case with an aldehyde, or having another carbon chain here, with the case of a ketone, we have an OH group. And you're probably saying, hey, why is this called an acid? It must be called an acid because it is an acid, and you'd be right. And the reason why it is an acid, and why it is more acidic that just something with an OH group, it's actually a good bit more acidic than alcohols, is because once this thing loses its hydrogen-- remember, depending on what type of acid you want to think about, acid in the Lewis sense could be an electron taker, and this oxygen can take the electron of this hydrogen. Or if you think about the Arrhenius definition of an acid, it is a proton donor, and this OH group can donate a proton. It takes away the electron of this hydrogen, gives the way the proton, either way. But the reason why this is more acidic than alcohol is once it gives away this proton, it is actually resonance stabilized. So let me show you what that means. To do that, let me actually show you the bond between this oxygen and this hydrogen. The oxygen has this pink electron, and the hydrogen has that magenta electron right there. If you put this in a solution of water, so you have some H2O over here, this oxygen right here really wants to take back this magenta electron. And as it takes back that magenta electron, it would essentially donate the hydrogen proton to a water molecule. So the water molecule would give one of its electrons to the hydrogen proton and then become positive. So once this happens, the next step would look like this. What was our carboxylic acid will now turn into the carboxylate ion, so it will now look like this, so it has our carbonyl group. Now, this oxygen just took an extra electron. So if I want to draw it, I have that one. Actually, let me draw. So this is the oxygen. To start off with, the oxygen had two lone pairs, so I want to draw those two lone pairs first. So to start off with, it had those two lone pairs, and now it had this pink electron from the get go, and now it took this magenta electron, so now it has one extra valence electron. We can even draw them here. We can count them: one, two, three, four, five, six. Six is just a neutral oxygen, but now it gained another one. It has seven. It now has a negative charge, and then the water has now become a hydronium ion. So you have the water here. We've increased the proton or the hydronium concentration in the water. This one water molecule is now a hydronium molecule, so this is now bonded with this hydrogen proton just like this. This oxygen gave away an electron to this proton, so now it has a positive charge. And this right here, this carboxylate ion right over here, the reason why this thing was a stronger acid than something that just had an OH group is because the conjugate base, the carboxylate ion, is actually resonance stabilized. It is more stable than the conjugate base of an alcohol, and let me show you that. This thing can share its negative charge. Let me draw it. It can take this magenta electron, give it to this carbon, then this carbon will have an extra electron, so then it can give back an electron to this top oxygen. So it is resonance stabilized with this structure right over here. Let me draw the same so it could look like this. That's too big. Let me scroll down a little bit. It could look like this. And now, this took back this blue electron, so now one of the bonds is gone, and it started off with two lone pairs, so I want to draw that there, and now it has another lone pair. It has this electron, this electron, and now it has that blue electron over there. And now this oxygen, this top oxygen, has a negative charge, and now the carbon has a double bond with this side oxygen. So now the carbon-- let me go back to the yellow-- that's the first bond with that oxygen. It had one, two lone pairs to begin with, and now it has this magenta bond. So this pink electron is at this end. And now this purple electron is at the other end, or this magenta electron, and now it has a double bond with this oxygen. And we know that when you have resonance stabilization, it's not like you're going back and forth. The reality is that you have a half double bond between both oxygens, that the electrons are just flowing across the whole place, and that stabilizes the molecule. And so to show that this is a resonance structure, let me put some brackets around it. And, in general, if this R group right here is actually even better at withdrawing electrons, so if you put something that was really electronegative here, something that likes to hog electrons, it would make the carboxylate ion even more stable. It would make the carboxylic acid even a better acid. So if you put something electronegative here, then you could imagine that some of this negative charge that we drew in these two resonance structures, can be sucked to that R group, and then that would make it even more stable, and would make the carboxylic acid even more acidic. Now, like in everything we've looked at, there are some common carboxylic acids that are not systematically named, that it's probably a good idea to know. And I'll start with one, just to see the pattern that we've seen in other things. We've seen that this thing over here, if we have this, we call this acetaldehyde. This was the very simple aldehyde we studied. we saw that if we have something like this, which is a ketone, we called this acetone. So you could imagine what we're probably going to name-- let me do this in a different color-- this molecule right over here. This is a carboxylic acid, clearly, and it has just that one methyl group, just like the acetaldehyde, just like the acetone, and so this is acetic acid. So the acetaldehyde, acetone, and acetic acid for me are fairly easy to memorize just because they all have the acet- part. They all have that as their prefix, and this part of the molecule is identical in every case. And the difference is the hydrogen, the methyl group over here, or the OH group, making this one a carboxylic acid, an acetone, and all that. Now, a couple of other ones that it wouldn't hurt for you to know is this one right over here. This is, you could argue, even simpler than acetic acid, and this formic acid. And then another one, and actually, I recently did a general chemistry video with this, where we did a titration example, but this is essentially two carboxyl groups attached to each other. It looks like this, two carboxyl groups attached to each other and you can see them. This is one carboxyl group right over here, and then you have another carboxyl group right over here. And so this actually can be deprotonated twice. This hydrogen can be lost, and that hydrogen can be lost, and this is oxalic acid. I'll leave you there. And in the next video, we'll learn how to systematically name carboxylic acids.