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.