Van Der Waals Forces

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

Throughout our journey through chemistry so farŁŹ

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 about it's good to have a general discussion

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 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 has 1 2 3 4 5 6 7 8...

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 haphazar 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 chlorin eatoms 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,

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,

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
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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.