Solubility and intermolecular forces | AP Chemistry | Khan Academy

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. Just as an example, we could go to our old friend sodium chloride and think about why it dissolves well in water.

Well, to do that, you just have to remind yourself what water is doing when nothing is dissolved in it. When nothing is dissolved in it, that's an oxygen attached to two hydrogens. This end, we've talked about in other videos, is partially negative because the electrons like to spend more time around the oxygen. This is partially positive; partially positive.

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. 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, can disassociate into its constituent ions: into a sodium cation and a chloride anion. We've seen this before.

So that's a 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 molecule, so the hydrogen end. 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 is stronger than the attraction between the ions themselves and the attractions between the water molecules themselves. 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, they have those weak forces kind of attracting them to each other. And if you think about a vat of hexane, 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 are many other anals, many other alcohols, but this is ethanol here. 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 is more electronegative than the things that it is bonded to. It still hogs electrons a little bit, and so you still have that partially negative charge; you still have that polarity to the molecule.

That's able to attract it to neighboring water molecules, which allows it to dissolve. But you can imagine if you had an alcohol that had a much longer carbon chain—say 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 will become, it'll make it harder and harder for it to dissolve in a polar solvent like water.

And so a mega takeaway here is that like dissolves like.