 OK. So this question asks, well, it asks a few questions. But the first of the questions that it asks are, which of these molecules exhibit London forces? Do you remember London forces? Were those very weak forces that there were the only forces that non-polar molecules have? So first thing I guess we've got to ask ourselves, what's polar and what's non-polar? And remember, hydrogen bonding. Remember, hydrogen bonding only happens when there's an H attached to a what? O or what else? And or an F. OK. Do we have something like that in here? Yeah, we do, right? We've got this guy over here, a methanol. So we've got an O attached to an H there. So we already know that that one's going to exhibit hydrogen bonding, OK? So this thing exhibits hydrogen bonding. I know I said London forces first, but we'll talk about it. That one's pretty obvious. Do the other ones have OH bonds or NH bonds or NF bonds? So none of them, none of the other ones will do hydrogen bonding, OK? So I'm going to ask those, which of these would exhibit dipole-dipole forces? Well, dipole-dipole forces, do you remember how I told you to remember how to do those ones? You looked at the table of electronegativity values, OK? And if they were different in the electronegativity values, then you had a dipole, OK? Remember I showed you with the little arrow? So if we look here, remember, we already know that oxygen and hydrogen have that dipole, right? So there's going to be a dipole-dipole force here with the carbon oxygen. And the hydrogen bonds are super dipole-dipole, OK? So if you're going to talk directly just dipole-dipole, you can see it with the carbon and oxygen. So this is definitely dipole-dipole, OK? And now let's look at this one here. Chlorine and carbon. Do they have a difference in electronegativity? Not by much, but do you guys have an electronegativity table, does anybody? OK, so mine says that there are three is chlorine and carbon is 2.5. So that's fairly close. When you get to about 0.4, the bias is the units for this. That's essentially non-polar, OK? So this one is very slightly polar, this bond here, so that you're going to have a dipole there, too. Dipole-dipole. If they asked you to show the dipole arrow, the one with the bigger electronegativity is more or stronger. Remember the tug-of-war thing we were talking about? So this one would look like that, OK? That dipole arrow. Or you could draw it like this. Partial negative, partial positive. Oh, you made it. We're all waiting for it, OK? So if we look at carbon and hydrogen, what's the difference between those two? And the electronegativity table, those of you who have it. Sure. 0.4, OK? So non-polar, OK? So anything with just hydrogen and carbon bonds is non-polar, OK? So do non-polar things have dipole-dipoles? Not non-polar. They only have, what? London forces, remember? I told you that, OK? So this thing has, what then? London forces, OK? Does this thing have London forces? How do we know? This is the total level. Any molecule has London forces, OK? So that's why I went back. And it's the only thing that non-polars have, but all the other molecules have them, too, OK? Because they have, remember the London forces were just the instantaneous dipole and the induced dipole. So it's all about, do you have electrons? Does this molecule have electrons? Yes, of course. Is this molecule? Yes, so they can move around and induce and have an instantaneous. So does this one have London's then? Yes. And then does this one? Yes. OK, so are we cool with figuring that out? If we wanted to do dipole arrow here, we could do this one, but it's too dipole-arrow. This portion of this molecule is polar. This portion of the molecule is non-polar. So this is like a function of a molecule for some reason. The last question that it asks is, which of these molecules would you expect to have the highest boiling point? Which one? One, two, or three? Left, middle, or up? The highest boiling point. Three, why? Remember the vacuum example that I