London dispersion forces | Intermolecular forces and properties | AP Chemistry | Khan Academy

What we're going to do in this video is start talking about forces that exist between even neutral atoms or neutral molecules. The first of these intermolecular forces we will talk about are London dispersion forces. So it sounds very fancy, but it's actually a pretty interesting and almost intuitive phenomenon.

We are used to thinking about atoms unless they say we have a neutral atom. So it has the same number of protons and electrons. And so that's all, those are all the protons and the neutrons in the nucleus, and then it'll have a cloud of electrons. So I'm just imagining all these electrons kind of jumping around; that's how I'm going to represent it.

Let's imagine, and this is definitely not drawn to scale; the nucleus would actually be much smaller if it was. But let's say that there is an adjacent atom right over here, and it's also neutral. Maybe it's the same type of atom; it could be different, but we're going to say it's neutral and it also has an electron cloud.

So if these are both neutral in charge, how would they be attracted to each other? And that's what London dispersion forces act to explain. Because we have observed that even neutral atoms and neutral molecules can get attracted to each other.

The way to think about it is electrons are constantly jumping around probabilistically. They're in this probability density cloud where an electron could be anywhere at any given moment, but they're not always going to be evenly distributed. You can imagine that there is a moment where that left atom might look like this, just for a moment, where most of, or maybe slightly more of, the electrons are spending time on the left side of the atom than on the right side.

So maybe it looks something like that. And so for that brief moment, you have a partial negative charge. This is the Greek letter delta, lower case delta, which is used to denote partial charge. On this side, you might have a partial positive charge. Because remember, when it was evenly distributed, the negative charge was offset by the positive charge of the nucleus.

But here, on the right side, because there's fewer electrons here, maybe you have a partial positive. On the left side, you have where most of the electrons are in that moment; partial negative. Now, what might this induce in the neighboring atom? Think about that. Pause the video. Think about what might happen in the neighboring atom.

Well, we know that like charges repel each other and opposite charges attract each other. So if we have a partial positive charge out here on the right side of this left atom, then the negative electrons might be attracted to it in this right atom. So these electrons here might actually be pulled a little bit to the left.

So they might be pulled a little bit to the left, and so that will induce what is called a dipole. Now you'll have a partial negative charge on the left side of this atom and then a partial positive charge on the right side of it. We already had a randomly occurring dipole on the left-hand side, but then that would have induced a dipole on the right-hand side.

A dipole is just when you have the separation of charge, where you have your positive and negative charges at two different parts of a molecule, or an atom, or really anything. But in this world, then all of a sudden these two characters are going to be attracted to each other, or the atoms are going to be attracted to each other.

This attraction that happens due to induced dipoles, that is exactly what London dispersion forces is all about. You can actually call London dispersion forces as induced dipole-induced dipole forces. They become attracted to each other because of what could start out as a temporary imbalance of electrons, but then it induces a dipole in the other atom or the other molecule, and then they get attracted.

So the next question you might ask is how strong can these forces get? And that's all about a notion of polarizability. How easy is it to polarize an atom or molecule? Generally speaking, the more electrons you have, so the larger the electron cloud, larger electron cloud, which is usually associated with molar mass. So usually molar mass, then the higher polarizability you're going to have because you're going to have more electrons to play around with.

If this was a helium atom, which has a relatively small electron cloud, you couldn't have a significant imbalance. At most, you might have two electrons on one side, which would cause some imbalance. But on the other hand, imagine a much larger atom or a much larger molecule. You could have much more significant imbalances—three, four, five, fifty electrons—and that would create a stronger temporary dipole, which would then induce a stronger dipole in the neighbors that could domino through the entire sample of that molecule.

For example, if you were to compare some noble gases to each other, and so we can look at the noble gases here on the right-hand side. If you were to compare the London dispersion forces between, say, helium and argon, which one would you think would have higher London dispersion forces? A bunch of helium atoms next to each other, or a bunch of argon atoms next to each other?

Well, the argon atoms have a larger electron cloud, so they have higher polarizability, and so you're going to have higher London dispersion forces. You can actually see that in their boiling points. For example, the boiling point of helium is quite low; it is negative 268.9 degrees Celsius, while the boiling point of argon, it's still at a low temperature by our standards, but it's a much higher temperature than the boiling point for helium. It's at negative 185.8 degrees Celsius.

So one way to think about this: if you were at, say, negative 270 degrees Celsius, you would find a sample of helium in a liquid state. But as you warm things up, as you get beyond negative 268.9 degrees Celsius, you're going to see that those London dispersion forces that are keeping those helium atoms together, sliding past each other in a liquid state, they're going to be overcome by the energy due to the temperature.

So they're going to be able to break free of each other, and essentially the helium is going to boil, and you're going to enter into a gaseous state—the state that most of us are used to seeing helium in. But that doesn't happen for argon until a good bit warmer, still cold by our standards. And that's because it takes more energy to overcome the London dispersion forces of argon because the argon atoms have larger electron clouds.

So generally speaking, the larger the molecule, because it has a larger electron cloud, it'll have higher polarizability and higher London dispersion forces. But also, the shape of the molecule matters. The more that the molecules can come in contact with each other, the more surface area they have exposed to each other, the more likely that they can induce these dipoles in each other.

For example, butane can come in two different forms. It can come in what's known as n-butane, which looks like this, so you have four carbons and ten hydrogens. So two, three, four, five, six, seven, eight, nine, ten. This is known as n-butane. But another form of butane, known as isobutane, would look like this.

So you have three carbons in the main chain, and then you have one carbon that breaks off of that middle carbon, and then they're all, they all have four bonds, and the leftover bonds you could say are with the hydrogens. So it would look like this: this right over here is isobutane.

Now, if you had a sample of a bunch of n-butane versus a sample of a bunch of isobutane, which of these do you think will have a higher boiling point? Pause this video and think about that. Well, if you have a bunch of n-butanes next to each other, imagine another n-butane right over here. It's going to have more surface area to its neighboring butanes because it is a long molecule.

It can expose that surface area to its neighbors, while the isobutane, in some ways, is a little bit more compact. It has lower surface area; it doesn't have these big long chains. And so because you have these longer n-butane molecules, you're going to have higher London dispersion forces. They obviously have the same number of atoms in them. They have the same number of electrons in them, so they have similar size electron clouds; they have the same molar mass.

But because of n-butane's elongated shape, they're able to get closer to each other and induce more of these dipoles. So just by looking at the shape of n-butane versus isobutane, you say higher London dispersion forces in n-butane, so it's going to have a higher boiling point. It's going to require more energy to overcome the London dispersion forces and get into a gaseous state.