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Throughout our journey through
chemistry so far, we've

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touched on the interactions
between molecules, metal

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molecules, how they attract each
other because of the sea

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of electrons and water
molecules.

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

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of the different types of
molecular interactions and

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what it means for the boiling
points or the melting points

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of a substance.

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

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bunch of helium.

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

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

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

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Because the point is that
noble gases are happy.

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Their outer orbital is filled.

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

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

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write neon like neon and
it's completely happy.

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It's completely satisfied
with itself.

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

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

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why it should be-- if these
electrons are evenly

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distributed around these
atoms, then these are

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

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

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

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or not attracted
to each other.

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

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

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

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

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

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

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it'll be a gas at most
temperatures.

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

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that starts to connect or makes
the neon molecules want

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to get towards each other.

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And that force comes out of
the reality that we talked

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

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orbit around things.

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They're probablistic.

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

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these nice and neat valence
dot electrons like that,

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

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probability cloud and it's
what neon's atomic

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configuration is.

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

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

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look-- I don't know.

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It has the 2s shell.

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

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The p-orbitals look like that
in different dimensions.

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That's not the point.

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

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just drawing the probability
distribution.

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I'm not trying to
draw a rabbit.

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But I think you get the point.

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Watch the electron configuration
videos if you

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want more on this, but the idea
behind these probability

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distributions is that the
electrons could be anywhere.

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There could be a moment in time
when all the electrons

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are out over here.

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There could be a moment in time
where all the electrons

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are over here.

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Same thing for this neon atom.

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

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

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

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to be completely evenly
distributed.

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There's many more scenarios
where the electron

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distribution is a little
uneven in one

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neon atom or another.

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So if in this neon atom,
temporarily its eight valence

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

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

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neon atom look like?

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It temporarily has a
slight charge in

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this direction, right?

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

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or this side is more positive
than that side.

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

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

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

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

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four, five, six, seven, eight.

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

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faint force.

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So this would be a
little negative.

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Temporarly, just for that single
moment in time, this

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will be kind of negative.

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That'll be positive.

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This side will be negative.

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This side will be positive.

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

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for that very small moment of
time between this neon and

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

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electrons will reconfigure.

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But the important thing to
realize is that almost at no

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point is neon's electrons
going to be completely

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distributed.

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

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

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

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almost too strong of a word.

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

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

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it to attract it to the opposite
side charges of other

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similarly imbalanced
molecules.

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

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It's called the London
dispersion force.

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

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was neither-- well, he
was not British.

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I think he was German-American.

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

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der Waals forces.

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

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

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intermolecular, and in
this case, neon-- the

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molecule, is an atom .

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

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

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intermolecular forces that are
not covalent bonds and that

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

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on those in a second.

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And the weakest of them are the
London dispersion forces.

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

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gases right here, the only thing
that they experience are

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London dispersion forces, which
are the weakest of all

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of the intermolecular forces.

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

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into a gaseous state.

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

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will turn into the
gaseous state.

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

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

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because they have
very, very small

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attraction to each other.

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Fair enough.

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

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molecules that have better
attractions or that are a

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little bit more polar?

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Let's say I had hydrogen
chloride, right?

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

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not it keeps its electrons.

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Chloride wants to keep
the electrons.

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Chloride's quite
electronegative.

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It's less electronegative than
these guys right here.

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These are kind of the
super-duper electron hogs,

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nitrogen, oxygen, and fluorine,
but chlorine is

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pretty electronegative.

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So if I have hydrogen chloride,
so I have the

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chlorine atom right here, it has
seven electrons and then

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it shares an electron
with the hydrogen.

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

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just do it like that.

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Because this is a good bit
more electronegative than

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

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

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side, where the electron
hog is, and a

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partial positive side.

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And this is actually
very analogous to

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the hydrogen bonds.

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

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

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

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another chlorine atom,
the other chlorine

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atoms looks like this.

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

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

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attraction between them.

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You'll have this attraction
between these two chlorine

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atoms-- oh, sorry,
between these two

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hydrogen chloride molecules.

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And the positive side, the
positive pole of this dipole

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is the hydrogen side, because
the electrons have kind of

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left it, will be attracted
to the chlorine side

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of the other molecules.

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

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interaction is stronger than
a London dispersion force.

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

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molecular interactions.

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

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much anything else.

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It only becomes relevant when
you talk about things with

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noble gases.

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

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electron distribution just
happens to go one way or the

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other for a single
instant of time.

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But this dipole-dipole
interaction is much stronger.

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And because it's much stronger,
hydrogen chloride is

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

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

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just a sample of helium gas.

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Now, when you get even more
electronegative, when this

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guy's even more electronegative
when you're

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

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

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hydrogen bond.

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

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a bunch of hydrogen fluorides
around the place.

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

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fluoride here.

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Fluoride its
ultra-electronegative.

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

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Periodic Table, and
so it pretty much

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hogs all of the electrons.

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

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

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be hogged around the
fluorine side.

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

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partial negative side, partial
positive, partial negative,

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partial positive, partial
negative and so on.

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

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interaction.

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

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

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very electronegative atom, where
the electronegative atom

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

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

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

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attracted to the negative
side of these molecules.

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But hydrogen, all of these
are van der Waals.

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So van der Waals, the weakest
is London dispersion.

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Then if you have a molecule with
a more electronegative

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

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side where molecule becomes
polar and you have the

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interaction between the positive
and the negative side

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of the pole.

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It gets a dipole-dipole
interaction.

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

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because the
super-electronegative atom is

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essentially stripping off the
electron of the hydrogen, or

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almost stripping it off.

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

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that side of the molecule.

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Since this is even a stronger
bond between molecules, it

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will have even a higher
boiling point.

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

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and then you have
hydrogen bonds.

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All of these are van der
Waals, but because the

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00:11:09,310 --> 00:11:13,360
strength of the intermolecular
bond gets stronger, boiling

245
00:11:13,360 --> 00:11:18,300
point goes up because it takes
more and more energy to

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00:11:18,300 --> 00:11:21,390
separate these from
each other.

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

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

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types of intermolecular
interactions that aren't

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00:11:28,370 --> 00:11:29,950
necessarily covalent or ionic.

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

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ionic types of structures that
can be formed and how that

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might affect the different
boiling points.