Van Der Waals Forces

Van Der Waals Forces: London Dispersion Forces, Dipole Attractions, and Hydrogen Bonds.

Throughout our journey through
chemistry so far, we've

touched on the interactions
between molecules, metal

molecules, how they attract each
other because of the sea

of electrons and water
molecules.

But I think it's good to have a
general discussion about all

of the different types of
molecular interactions and

what it means for the boiling
points or the melting points

of a substance.

So I'll start with the weakest.
Let's say I had a

bunch of helium.

Helium, you know, I'll just draw
it as helium atoms. We'll

look in the Periodic Table, and
what I'm going to do now

with helium I could do with
any of the noble gases.

Because the point is that
noble gases are happy.

Their outer orbital is filled.

Let's say, neon or helium--
let me do neon, actually,

because neon has a full eight
in its orbital so we could

write neon like neon and
it's completely happy.

It's completely satisfied
with itself.

And so in a world where it's
completely satisfied, there's

no obvious reason just yet-- I'm
going to touch on a reason

why it should be-- if these
electrons are evenly

distributed around these
atoms, then these are

completely neutral atoms. They
don't want to bond with each

other or do anything else, so
they should just float around

and there's no reason for them
to be attracted to each other

or not attracted
to each other.

But it turns out that neon does
have a liquid state, if

you get cold enough, and so the
fact that it has a liquid

state means that there must be
some force that's making the

neon atoms attracted to each
other, some force out there.

Because it's in a very cold
state, because for the most

part, there is not a lot of
force that attracts them so

it'll be a gas at most
temperatures.

But if you get really cold, you
can get a very weak force

that starts to connect or makes
the neon molecules want

to get towards each other.

And that force comes out of
the reality that we talked

about early on that electrons
are not in a fixed, uniform

orbit around things.

They're probablistic.

And if we imagine, let me say
neon now, instead of drawing

these nice and neat valence
dot electrons like that,

instead, I can kind of draw
its electrons as-- it's a

probability cloud and it's
what neon's atomic

configuration is.

1s2 and it's outer orbital
is 2s2 2p6, right?

So it's highest energy electron,
so, you know, it'll

look-- I don't know.

It has the 2s shell.

The 1s shell is inside of that
and it has the p-orbitals.

The p-orbitals look like that
in different dimensions.

That's not the point.

And then you have another neon
atom and these are-- and I'm

just drawing the probability
distribution.

I'm not trying to
draw a rabbit.

But I think you get the point.

Watch the electron configuration
videos if you

want more on this, but the idea
behind these probability

distributions is that the
electrons could be anywhere.

There could be a moment in time
when all the electrons

are out over here.

There could be a moment in time
where all the electrons

are over here.

Same thing for this neon atom.

If you think about it, out
of all of the possible

configurations, let's say we
have these two neon atoms,

there's actually a very low
likelihood that they're going

to be completely evenly
distributed.

There's many more scenarios
where the electron

distribution is a little
uneven in one

neon atom or another.

So if in this neon atom,
temporarily its eight valence

electrons just happen to be
like, you know, one, two,

three, four, five, six, seven,
eight, then what does this

neon atom look like?

It temporarily has a
slight charge in

this direction, right?

It'll feel like this side is
more negative than this side

or this side is more positive
than that side.

Similarly, if at that very same
moment I had another neon

that had one, two, three, four,
five, six, seven, eight,

that had a similar-- actually,
let me do that differently.

Let's say that this neon atom is
like this: one, two, three,

four, five, six, seven, eight.

So here, and I'll do it in a
dark color because it's a very

faint force.

So this would be a
little negative.

Temporarly, just for that single
moment in time, this

will be kind of negative.

That'll be positive.

This side will be negative.

This side will be positive.

So you're going to have a little
bit of an attraction

for that very small moment of
time between this neon and

this neon, and then it'll
disappear, because the

electrons will reconfigure.

But the important thing to
realize is that almost at no

point is neon's electrons
going to be completely

distributed.

So as long as there's always
going to be this haphazard

distribution, there's always
going to be a little bit of

a-- I don't want to say polar
behavior, because that's

almost too strong of a word.

But there will always be a
little bit of an extra charge

on one side or the other side
of an atom, which will allow

it to attract it to the opposite
side charges of other

similarly imbalanced
molecules.

And this is a very, very,
very weak force.

It's called the London
dispersion force.

I think the guy who came up with
this, Fritz London, who

was neither-- well, he
was not British.

I think he was German-American.

London dispersion force, and
it's the weakest of the van

der Waals forces.

I'm sure I'm not pronouncing
it correctly.

And the van der Waals forces
are the class of all of the

intermolecular, and in
this case, neon-- the

molecule, is an atom .

It's just a one-atom molecule,
I guess you could say.

The van der Waals forces are
the class of all of the

intermolecular forces that are
not covalent bonds and that

aren't ionic bonds like we have
in salts, and we'll touch

on those in a second.

And the weakest of them are the
London dispersion forces.

So neon, these noble gases,
actually, all of these noble

gases right here, the only thing
that they experience are

London dispersion forces, which
are the weakest of all

of the intermolecular forces.

And because of that, it takes
very little energy to get them

into a gaseous state.

So at a very, very low
temperature, the noble gases

will turn into the
gaseous state.

That's why they're called noble
gases, first of all.

And they're the most likely
to behave like ideal gases

because they have
very, very small

attraction to each other.

Fair enough.

Now, what happens when we go
to situations when we go to

molecules that have better
attractions or that are a

little bit more polar?

Let's say I had hydrogen
chloride, right?

Hydrogen, it's a little bit
ambivalent about whether or

not it keeps its electrons.

Chloride wants to keep
the electrons.

Chloride's quite
electronegative.

It's less electronegative than
these guys right here.

These are kind of the
super-duper electron hogs,

nitrogen, oxygen, and fluorine,
but chlorine is

pretty electronegative.

So if I have hydrogen chloride,
so I have the

chlorine atom right here, it has
seven electrons and then

it shares an electron
with the hydrogen.

It shares an electron with
the hydrogen, and I'll

just do it like that.

Because this is a good bit
more electronegative than

hydrogen, the electrons spend
a lot of time out here.

So what you end up having is a
partial negative charge on the

side, where the electron
hog is, and a

partial positive side.

And this is actually
very analogous to

the hydrogen bonds.

Hydrogen bonds are actually a
class of this type of bond,

which is called a dipole bond,
or dipole-dipole interaction.

So if I have one chlorine atom
like that and if I have

another chlorine atom,
the other chlorine

atoms looks like this.

If I have the other chlorine
atom-- let me copy and paste

it-- right there, then
you'll have this

attraction between them.

You'll have this attraction
between these two chlorine

atoms-- oh, sorry,
between these two

hydrogen chloride molecules.

And the positive side, the
positive pole of this dipole

is the hydrogen side, because
the electrons have kind of

left it, will be attracted
to the chlorine side

of the other molecules.

And because this van der Waals
force, this dipole-dipole

interaction is stronger than
a London dispersion force.

And just to be clear, London
dispersion forces occur in all

molecular interactions.

It's just that it's very weak
when you compare it to pretty

much anything else.

It only becomes relevant when
you talk about things with

noble gases.

Even here, they're also London
dispersion forces when the

electron distribution just
happens to go one way or the

other for a single
instant of time.

But this dipole-dipole
interaction is much stronger.

And because it's much stronger,
hydrogen chloride is

going to take more energy to,
one, get into the liquid

state, or even more, get into
the gaseous state than, say,

just a sample of helium gas.

Now, when you get even more
electronegative, when this

guy's even more electronegative
when you're

dealing with nitrogen, oxygen
or fluorine, you get into a

special case of dipole-dipole
interactions, and that's the

hydrogen bond.

So it's really the same thing if
you have hydrogen fluoride,

a bunch of hydrogen fluorides
around the place.

Maybe I could write fluoride,
and I'll write hydrogen

fluoride here.

Fluoride its
ultra-electronegative.

It's one of the three most
electronegative atoms on the

Periodic Table, and
so it pretty much

hogs all of the electrons.

So this is a super-strong case
of the dipole-dipole

interaction, where here, all of
the electrons are going to

be hogged around the
fluorine side.

So you're going to have a
partial positive charge,

partial negative side, partial
positive, partial negative,

partial positive, partial
negative and so on.

So you're going to have this,
which is really a dipole

interaction.

But it's a very strong dipole
interaction, so people call it

a hydrogen bond because it's
dealing with hydrogen and a

very electronegative atom, where
the electronegative atom

is pretty much hogging all of
hydrogen's one electron.

So hydrogen is sitting out here
with just a proton, so

it's going to be pretty
positive, and it's really

attracted to the negative
side of these molecules.

But hydrogen, all of these
are van der Waals.

So van der Waals, the weakest
is London dispersion.

Then if you have a molecule with
a more electronegative

atom, then you start having a
dipole, where you have one

side where molecule becomes
polar and you have the

interaction between the positive
and the negative side

of the pole.

It gets a dipole-dipole
interaction.

And then an even stronger type
of bond is a hydrogen bond

because the
super-electronegative atom is

essentially stripping off the
electron of the hydrogen, or

almost stripping it off.

It's still shared,
but it's all on

that side of the molecule.

Since this is even a stronger
bond between molecules, it

will have even a higher
boiling point.

So London dispersion, and you
have dipole or polar bonds,

and then you have
hydrogen bonds.

All of these are van der
Waals, but because the

strength of the intermolecular
bond gets stronger, boiling

point goes up because it takes
more and more energy to

separate these from
each other.

In the next video-- I realize
I'm out of time.

So this is a good survey, I
think, of just the different

types of intermolecular
interactions that aren't

necessarily covalent or ionic.

In the next video, I'll talk
about some of the covalent and

ionic types of structures that
can be formed and how that

might affect the different
boiling points.