is start talking about forces

that exist between even neutral atoms

or neutral molecules.

And 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.

It has the same number of protons and electrons.

And so those are all the protons

and the neutrons and 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 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 gonna say it's neutral.

And it also has an electron cloud.

And so if these are both neutral and charged,

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 the electron could be anywhere at any given moment

but they're not always going to be evenly distributed.

You can imagine that there's a moment

where that left atom might look like this,

just for a 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,

and 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 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

side but then that would've 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

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 induced dipole, induced dipole forces.

They become attracted to each other

because of what could start out as 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, electron cloud,

which is usually associated with molar mass,

so usually molar mass, then

the higher polarizability you're gonna have,

'cause 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, 50 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.

So 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

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 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 both a sample of helium

and a sample of argon in a liquid state.

They would each be in liquid states,

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

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'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 10 hydrogens, so two,

three, four, five, six, seven, eight, nine, 10.

This is known as n-butane,

but another form of butane known as iso

butane would look like this.

So you have 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 leftover bonds you could say

are with the hydrogens.

So it would look like this.

This right over here is iso butane.

Iso butane.

Now, if 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 other 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.

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-butanes' elongated shape,

they are 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 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.