In the video on
electronegativity, we learned how to determine
whether a covalent bond is polar or nonpolar. In this video, we're going
to see how we figure out whether molecules
are polar or nonpolar and also how to apply
that polarity to what we call intermolecular forces. Intermolecular
forces are the forces that are between molecules. And so that's different from
an intramolecular force, which is the force within a molecule. So a force within
a molecule would be something like
the covalent bond. And an intermolecular
force would be the force that are
between molecules. And so let's look at the
first intermolecular force. It's called a
dipole-dipole interaction. And let's analyze
why it has that name. If I look at one of these
molecules of acetone here and I focus in on the
carbon that's double bonded to the oxygen,
I know that oxygen is more electronegative
than carbon. And so we have four
electrons in this double bond between the carbon
and the oxygen. So I'll try to highlight
them right here. And since oxygen is
more electronegative, oxygen is going to pull
those electrons closer to it, therefore giving oxygen a
partial negative charge. Those electrons in yellow are
moving away from this carbon. So the carbon's losing a
little bit of electron density, and this carbon is becoming
partially positive like that. And so for this
molecule, we're going to get a separation of charge, a
positive and a negative charge. So we have a polarized
double bond situation here. We also have a
polarized molecule. And so there's two
different poles, a negative and a positive pole here. And so we say that this
is a polar molecule. So acetone is a
relatively polar molecule. The same thing happens to this
acetone molecule down here. So we get a partial negative,
and we get a partial positive. So this is a polar
molecule as well. It has two poles. So we call this a dipole. So each molecule
has a dipole moment. And because each
molecule is polar and has a separation of
positive and negative charge, in organic chemistry we know
that opposite charges attract, right? So this negatively
charged oxygen is going to be attracted to
this positively charged carbon. And so there's going to be
an electrostatic attraction between those two molecules. And that's what's going to hold
these two molecules together. And you would
therefore need energy if you were to try
to pull them apart. And so the boiling
point of acetone turns out to be approximately
56 degrees Celsius. And since room temperature
is between 20 and 25, at room temperature
we have not reached the boiling point of acetone. And therefore, acetone
is still a liquid. So at room temperature and
pressure, acetone is a liquid. And it has to do with
the intermolecular force of dipole-dipole
interactions holding those
molecules together. And the intermolecular
force, in turn, depends on the
electronegativity. Let's look at another
intermolecular force, and this one's called
hydrogen bonding. So here we have two
water molecules. And once again, if I think
about these electrons here, which are between the
oxygen and the hydrogen, I know oxygen's more
electronegative than hydrogen. So oxygen's going to pull
those electrons closer to it, giving the oxygen a partial
negative charge like that. The hydrogen is losing a
little bit of electron density, therefore becoming
partially positive. The same situation exists in
the water molecule down here. So we have a partial negative,
and we have a partial positive. And so like the
last example, we can see there's going
to be some sort of electrostatic attraction
between those opposite charges, between the negatively
partially charged oxygen, and the partially positive
hydrogen like that. And so this is a polar molecule. Of course, water is
a polar molecule. And so you would
think that this would be an example of
dipole-dipole interaction. And it is, except
in this case it's an even stronger version of
dipole-dipole interaction that we call hydrogen bonding. So at one time it
was thought that it was possible for hydrogen
to form an extra bond. And that's where the term
originally comes from. But of course, it's not an
actual intramolecular force. We're talking about an
intermolecular force. But it is the strongest
intermolecular force. The way to recognize when
hydrogen bonding is present as opposed to just
dipole-dipole is to see what the hydrogen is bonded to. And so in this case, we have
a very electronegative atom, hydrogen, bonded-- oxygen,
I should say-- bonded to hydrogen. And then that hydrogen
is interacting with another electronegative
atom like that. So we have a partial negative,
and we have a partial positive, and then we have another
partial negative over here. And this is the
situation that you need to have when you
have hydrogen bonding. Here's your hydrogen showing
intermolecular force here. And what some students forget
is that this hydrogen actually has to be bonded to another
electronegative atom in order for there to be a big enough
difference in electronegativity for there to be a little
bit extra attraction. And so the three
electronegative elements that you should remember
for hydrogen bonding are fluorine,
oxygen, and nitrogen. And so the mnemonics
that students use is FON. So if you remember FON as the
electronegative atoms that can participate in
hydrogen bonding, you should be able to remember
this intermolecular force. The boiling point of water is,
of course, about 100 degrees Celsius, so higher than
what we saw for acetone. And this just is due to the
fact that hydrogen bonding is a stronger version of
dipole-dipole interaction, and therefore, it takes
more energy or more heat to pull these water
molecules apart in order to turn
them into a gas. And so, of course, water is
a liquid at room temperature. All right. Let's look at another
intermolecular force. And this one is called
London dispersion forces. So these are the weakest
intermolecular forces, and they have to do with the
electrons that are always moving around in orbitals. And even though the
methane molecule here, if we look at it,
we have a carbon surrounded by four
hydrogens for methane. And it's hard to tell in how
I've drawn the structure here, but if you go back and
you look at the video for the tetrahedral
bond angle proof, you can see that in
three dimensions, these hydrogens are
coming off of the carbon, and they're equivalent
in all directions. And there's a very
small difference in electronegativity between
the carbon and the hydrogen. And that small difference
is canceled out in three dimensions. So the methane molecule becomes
nonpolar as a result of that. So this one's nonpolar, and,
of course, this one's nonpolar. And so there's no
dipole-dipole interaction. There's no hydrogen bonding. The only intermolecular
force that's holding two methane
molecules together would be London
dispersion forces. And so once again, you could
think about the electrons that are in these bonds
moving in those orbitals. And let's say for the
molecule on the left, if for a brief
transient moment in time you get a little bit
of negative charge on this side of the molecule,
so it might turn out to be those electrons have a net
negative charge on this side. And then for this
molecule, the electrons could be moving the
opposite direction, giving this a partial positive. And so there could be
a very, very small bit of attraction between these
two methane molecules. It's very weak, which is why
London dispersion forces are the weakest
intermolecular forces. But it is there. And that's the only thing that's
holding together these methane molecules. And since it's weak, we would
expect the boiling point for methane to be extremely low. And, of course, it is. So the boiling point for methane
is somewhere around negative 164 degrees Celsius. And so since room temperature
is somewhere around 20 to 25, obviously methane
has already boiled, if you will, and
turned into a gas. So methane is obviously a gas at
room temperature and pressure. Now, if you increase
the number of carbons, you're going to increase the
number of attractive forces that are possible. And if you do that,
you can actually increase the boiling point
of other hydrocarbons dramatically. And so even though
London dispersion forces are the weakest, if you
have larger molecules and you sum up all
those extra forces, it can actually turn out to be
rather significant when you're working with larger molecules. And so this is just
a quick summary of some of the
intermolecular forces to show you the application
of electronegativity and how important it is.