Hi. It's Mr. Andersen and this is chemistry essentials video 16. It's on
London dispersion forces which were first identified by Fritz London. That's where they
get their name. After this whole idea of quantum mechanics had been developed. Can you find
him in this picture? He is right here. Before we move on and talk about it, think about
the forces in this picture. Which are greater? The forces within the individual people holding
their atoms together? Or the forces between them holding them together in this photo?
Well clearly it's the forces within your body. And so this is the first video in a set a
videos that talk about bonding between molecules and within molecules. And so if we think of
them all as molecules, and we move them apart a little bit, there are going to be the intermolecular
forces. Those are going to be the forces between the molecules. Between this molecule and this
molecule. And between this molecule and that molecule. And these are generally going to
be really, really low compared to the intra molecular, within the molecule itself. And
we'll talk more about those in later videos. But London dispersion forces are going to
be attractive forces. These are intermolecular, between molecules. And they're going to be
found in all atoms and molecules. Everything is attracted through London dispersion forces.
It's really important. What really develops them are uneven electrons. As the electrons
are distributed unevenly they create what are called a temporary dipole. Those are going
to be a molecule that has a charge. It has a positive and a negative charge. These are
going to increase as we increase the contact area between the molecules. Also as we increase
the polarizability, which is going to be the squashiness of a molecule. Generally, the
more electrons you have, the higher this is going to be. And so as we increase molecular
size we're going to increase these dispersion forces. And finally pi bonding, which is going
to be where orbitals overlap. If we have more pi bonding, we're going to increase these.
And so let's look at it. Why are London dispersion forces important? Well they're in everything.
And if it weren't for them we'd never have noble gases in a liquid phase. And we know
we can do that if we can cool it down. And so what holds helium together when it gets
into this liquid phase, is going to be these London dispersion forces. So right here I
have two helium atoms. And so they're going to have two protons on the inside. Two electrons
on the outside. And so as we move them close together, watch what happens to their electrons.
Let's watch that again. As they move close together the electrons are going to migrate.
What's doing that? Well the electrons on the outside of one helium are going to be attracted,
just Coulomb's law, to the protons on the inside of that other helium. There's also
going to be repulsion between those electrons on the outside. And so what this really is
creating are two, what are called temporary dipoles. They have positive and negative ends.
And so a London dispersion force is simply going to be the connection between those two.
It's fairly simple. It's going to hold them together. Now it's just instantaneous. It's
temporary. And then it's going to go away. But if we have enough of these atoms in an
area and it's cool enough, then we can get that. So what's going to increase these London
dispersion forces? Well contact area is one thing. The more electrons we can have closer
to each other, the more of these forces we're going to find. So if we look over here at
n-pentane & neopentane, they're both going to be made of the same atoms. But these ones
right here are going to be looking more like a column and these more like spheres. And
as a result, these ones are going to have more surface area, more contact area. So we're
going to have greater forces. And so the boiling point of n-pentane is going to be higher than
that of neopentane based on these London dispersion forces. What else? The polarizability. And
polarizability remember is the squashiness. The more electrons we have, the more forces
we're going to have. And so this right here is a protein. Proteins are massive in size.
And a lot of that 3-dimensional shape comes from these London dispersion forces. Also
it's going to explain why as we go down in the periodic table, let's look at the halogens
right here. Halogens are going to have the same valence electrons, but as we go to fluorine,
chlorine, bromine and iodine, we're actually going to go from gases, fluorine, chlorine,
to liquid bromine and then solid iodine. And that's all going to be based on these London
dispersion forces. As the atoms get bigger they have more electrons. We're going to have
more of this uneven electron. And then finally pi bonding. That's going to have orbitals that
overlap. It forms something called a pi bond. If we increase the number of pi bonds inside
those molecules, we're going to increase these London dispersion forces. So again, what are
they? They are attractive in nature. Found in everything. It's uneven electrons create
these temporary dipoles. Remember what increases it, our list would be contact area or surface
contact. Also the polarizability which is going to be based on the molecule size. And
pi bonding. Can you do that? Could you explain these trends? If you can remember those three
things, you've got it. And I hope that was helpful.