Nice explanation bro!
Nice explanation bro!
A small misspelling: fluorine has not positive (like on the video), but negative charge.
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.
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
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
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
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.
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
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
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
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
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 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
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.
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