- [Instructor] In this video, we're going to talk
about solubility,
which is just a way of describing how well certain solutes
can dissolve in certain solvents.
And just as an example, we could go to our old friend
sodium chloride and think about
why does it dissolve well in water.
Well, to do that, you just have to remind yourself
what water is doing when nothing is dissolved in it.
So when nothing is dissolved in it,
that's an oxygen attached to two hydrogens.
This end, we've talked about in other videos,
partially negative
'cause the electrons like to spend more time
around the oxygen.
This is partially positive, partially positive.
And you have this hydrogen bond forming,
where the partially negative end of one water molecule
is attracted to the partially positive end
of another water molecule.
That's what hydrogen bonds are all about.
And we have whole videos on that,
actually many videos on that.
The reason why sodium chloride dissolves well in water
is because sodium chloride, as an ionic compound,
it can disassociate into its constituent ions.
Into a sodium cation
and a chloride anion.
And we've seen this before.
So that's the chloride anion,
this right over here is a sodium cation.
The reason why this dissolves well
is that the negative charge is able to be drawn
to the positive end of the water molecules,
so the hydrogen end.
And the positive sodium cation is attracted
to the negative end of the water molecules.
So what's really happening is the attraction
between the ions and the water molecules are stronger
than the attraction between the ions themselves
and the attraction between the water molecules themselves.
So the water molecules, don't just bunch up and say,
"We want nothing to do with you, sodium and chloride."
They say, "Hey, we're kind of attracted to you too.
So why don't we mix together?"
We can look at things that have less attractive forces,
maybe things where the main force is just dispersion forces.
If you think about a vat of pentane here,
they have those weak forces,
kind of attracting them to each other,
and then if you think about a vat of hexane here,
there's kind of weak forces.
But if you were to put some pentane,
let's call this the solute here.
And if you were to put it into a solvent of hexane,
it will dissolve because they are roughly as attracted
to each other as they are to themselves.
Now, what do you think is going to happen
if I try to put, say, some hexane,
if I view that as a solute,
and I were to put it in water?
Well, in that situation,
the water is going to be far more attracted to itself
than it's going to be attracted to the hexane.
Let's say that this is the water here.
You're going to have these globs of the hydrocarbon form
because the water is more attracted to itself.
It's not easy for the hydrocarbon to dissolve.
Now, there are many organic molecules that do dissolve well.
And that's usually because they have some part
of the molecule that has some polarity to it.
One example is ethanol, which has an OH group.
But ethanol, which has a chain of two carbons,
when we talk about alcohol in everyday language,
drinking alcohol, that is ethanol.
There's many other anols, many other alcohols,
but this is ethanol here.
And if you were to take alcohol and you were to mix it
in water, it does dissolve well.
And that's because this oxygen here,
it's more electronegative
than the things that it is bonded to.
And so it still hogs electrons a little bit.
And so you still have that partially negative charge.
You still have that polarity to the molecule.
And so that's able to attract it
to neighboring water molecules, which allows it to dissolve.
But you could imagine if you had an alcohol
that had a much, much longer carbon chain,
so you had 10 carbons or 15 carbons,
then all of a sudden the relative proportion
of how polar it is compared to
how large of the molecule it is,
it'll make it harder and harder for it to dissolve
in a polar solvent like water.
And so our mega takeaway here is that like dissolves like.