arrange these compounds by their expected vapor pressure.

But then the things thatĪre in the gaseous state, every now and then they'reīumping into each other, and they're bumping into You're going to have a molecule that has the right positionĪnd the right kinetic energy to escape and get into the vapor state, into a gaseous state. The temperature, youĬould view as a measure of the average kineticĮnergy of the molecules, but they're all bumpingĪround into each other, in different positions, withĭifferent amounts of velocities and therefore different kinetic energies.

Of these molecules is moving with the exact same kinetic energy. Going to be in a liquid state, but we know that not every one Now, you might notice,Īt 20 degrees Celsius, it's lower than the boiling point of all of these characters. And the temperature matters, so let's say that this Molecules in a liquid state, and I'm gonna just draw the molecules, clearly not drawn to scale,Īs these little circles. And I put one of these, a sample of one of these But just to get you a sense, imagine a closed container here. So what are we talking about, why, about vapor pressure, and whyĭo we see this relationship? And I'm not going to goĭeep into vapor pressure. The low boiling point have a high vapor pressure. The high boiling point have the low vapor pressure, and the things that have That vapor pressure seems to trend the opposite Now, what's also interesting here, you might have noticed, is this We can see very clearly that water has the highest boiling point, ethanol is second, methanol is third, and diethyl ether was fourth, completely consistent with our intuition. And if we look at the actual data, it's consistent with what Won the tiebreaker, followed by ethanol, followed by methanol, and then the lowest boiling I would put would be water, followed by, since ethanol On boiling points, the highest boiling point But they can be useful for the tiebreaker between ethanol and methanol. And so London dispersion forces, I wouldn't make that change the ranking between water or diethylĮther because these are going to be a lot weaker than And you can literally take atoms away from that to get to a water. And you can literally take atoms away from that to get to a methanol. How did I know that? Well, you literally can take atoms away from the diethyl ether And it's clear that diethyl ether has the highest molar mass, followed by ethanol, followed by methanol, followed by water. Large its electron cloud is, which is proportional to its molar mass. So London dispersionįorces are proportional to how polarizable a molecule is, which is proportional to how And if we're just trying to, actually I'll rank all of them. Similar dipole moments on a molecular basis. Other types of dipole forces, but not a lot that you could But what about the differenceīetween methanol and ethanol? And we could think about So just looking at this, I know that water's going to And then I would put diethyl ether last 'cause it can't form hydrogen bonds. I would put methanol andĮthanol as a tie for second. Hydrogen bond contribution to the intermolecular forces, I would put water as number one 'cause it can form the We don't see any bonds between hydrogen and an oxygen, a nitrogen, or a fluorine. 'Cause you could really view those, those are the strongest of theĭipole-dipole interactions, and they're going to be stronger than your London dispersion forces. So I will start with hydrogen bonds, hydrogen bonds. Intermolecular forces that we have studied. Intermolecular forces and get to a gas state. Or a higher boiling point to really overcome those

Intermolecular forces, it would take a lot of energy Highest intermolecular forces when they're in a liquid state? Because if you have high All right, now to figure that out, it really just boils down to which of these has the

And what I want you to think about, if you had a pure sample of each, which of those pure samples would have the highest boiling point, second highest, third highest, and fourth highest? Pause this video, and

Arrange these compounds by their expected vapor pressure.