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.

So we are used to thinking about atoms,

and let's just say we have a neutral atom.

So it has the same number of proton and electrons.

And so 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

kinda jumping around.

That's how I'm going to represent it.

And 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's 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 gonna say it's neutral.

And it also has an electron cloud.

And so if these are both neutral in charge,

how would they be attracted to each other?

And that's what London dispersion forces actually explain.

Because we have observed that even neutral atoms

and neutral molecules can get attracted to each other.

And 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 moment,

where 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, lowercase delta,

which is used to denote partial charge.

And 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 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 and think about

what might happen in the neighboring atom then.

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,

well 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.

So 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.

And we already had a randomly occurring dipole

on the left hand side, but then that would have

induce 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.

And 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?

And 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 gonna have.

You're just gonna 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 neighbor.

That could domino through

the entire sample of that molecule.

So for example, if you were to compare

some noble gases to each other.

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 do 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.

And 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 for 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.

And 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 will 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.

Two, three, four, five, six,

seven, eight, nine, ten.

This is known as n-butane.

But another form of butane known as iso-butane

would look like this.

Three Carbons in the main chain,

then you have one Carbon that breaks off of that

middle Carbon and then they all have four bonds.

And the left over bonds, you could say,

are with the Hydrogens.

So it would look like this.

This right over here is iso-butane.

Now if you had a sample of a bunch of n-butane,

versus a sample of a bunch of iso-butane,

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 iso-butane 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 sized electron clouds.

The 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 iso-butane, you'd see higher

London dispersion forces in n-butane,

so its 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.