Van der waals forces - Chemistry, Class 11 Video Lecture

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.