- [Instructor] 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. 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.