B2 中高級 29 分類 收藏
- [Instructor] What we're going to do in this video
is start talking about forces
that exist between even neutral atoms
or neutral molecules.
And the first of these intermolecular forces
we will talk about are London dispersion forces.
So it sounds very fancy,
but it's actually a pretty interesting
and almost intuitive phenomenon.
So we are used to thinking about atoms.
And let's just say we have a neutral atom.
It has the same number of protons and electrons.
And so those are all the protons
and the neutrons and the nucleus.
And then it'll have a cloud of electrons.
So I'm just imagining all these electrons
kinda jumping around, that's how I'm going to represent it.
And let's imagine,
and this is definitely not drawn to scale,
the nucleus would actually be much smaller if it was,
but let's say that there is an adjacent atom,
right over here, and it's also neutral.
Maybe it's the same type of atom.
It could be different but we're gonna say it's neutral.
And it also has an electron cloud.
And so if these are both neutral and charged,
how would they be attracted to each other?
And that's what London dispersion forces actually explain,
because we have observed that
even neutral atoms and neutral molecules
can get attracted to each other.
And the way to think about it is,
electrons are constantly jumping around,
probabilistically, they're in this probability density cloud
where the electron could be anywhere at any given moment
but they're not always going to be evenly distributed.
You can imagine that there's a moment
where that left atom might look like this,
just for a moment,
where maybe slightly more of the electrons
are spending time on the left side of the atom
than on the right side.
So maybe it looks something like that.
And so for that brief moment,
you have a partial negative charge,
this is the Greek letter, delta, lowercase delta,
which is used to denote partial charge.
And on this side, you might have a partial positive charge.
Because remember, when it was evenly distributed,
the negative charge was offset
by the positive charge of the nucleus.
But here on the right side,
because there's fewer electrons here,
maybe you have a partial positive,
on the left side, where most of the electrons are,
and in that moment, partial negative.
Now what might this induce in the neighboring atom?
Think about that.
Pause the video.
Think about what might happen in the neighboring atom then?
Well, we know that like charges repel each other
and opposite charges attract each other.
So if we have a partial positive charge
out here on the right side of this left atom,
well then the negative electrons
might be attracted to it in this right atom.
So these electrons here might actually
be pulled a little bit to the left.
So they might be pulled a little bit to the left.
And so that will induce what is called a dipole.
So now you'll have a partial negative charge
on the left side of this atom,
and then a partial positive charge on the right side of it.
And we already had a randomly occurring dipole on the left
side but then that would've induced a dipole
on the right-hand side.
A dipole is just when you have the separation of charge
where you have your positive
and negative charges at two different parts of a molecule
or an atom or really anything.
But in this world, then
all of a sudden these two characters are
going to be attracted to each other,
or the atoms are going to be attracted to each other
and this attraction that happens due to induced dipoles,
that is exactly what London dispersion forces is all about.
You can actually call London dispersion
forces induced dipole, induced dipole forces.
They become attracted to each other
because of what could start out as temporary imbalance
of electrons, but then it induces a dipole in the other atom
or the other molecule and then they get attracted.
So the next question you might ask
is how strong can these forces get?
And that's all about a notion of polarizability,
how easy is it to polarize an atom or molecule,
and generally speaking,
the more electrons you have,
so the larger the electron cloud,
larger electron cloud, electron cloud,
which is usually associated with molar mass,
so usually molar mass, then
the higher polarizability you're gonna have,
'cause you're just gonna
have more electrons to play around with.
If this was a helium atom,
which has a relatively small electron cloud,
you couldn't have a significant imbalance.
At most you might have two electrons on one side,
which would cause some imbalance,
but on the other hand, imagine a much larger atom
or a much larger molecule.
You could have much more significant imbalances.
Three, four, five, 50 electrons,
and that would create a stronger temporary dipole
which would then induce a stronger dipole in the neighbors.
That could domino through
the entire sample of that molecule.
So for example, if you were
to compare some noble gases to each other,
and so we can look at the noble gases here
on the right-hand side,
if you were to compare the London dispersion
forces between say helium and argon,
which one would you think
have higher London dispersion forces?
A bunch of helium atoms next to each other
or a bunch of argon atoms next to each other?
Well, the argon atoms have a larger electron cloud,
so they have higher polarizability,
and so you're going to have higher London dispersion forces,
and you can actually see that in their boiling points.
For example, the boiling point of helium is quite low.
It is negative 268.9 degrees Celsius,
while the boiling point of argon,
it's still at a low temperature by our standards
but it's a much higher temperature
than the boiling point for helium.
It's at negative 185.8 degrees Celsius.
So one way to think about this,
if you were at say negative 270 degrees Celsius,
you would find both a sample of helium
and a sample of argon in a liquid state.
They would each be in liquid states,
but as you warm things up,
as you get beyond negative 268.9 degrees Celsius,
you're going to see that those London dispersion forces
are keeping those helium atoms together,
sliding past each other in a liquid state.
They're going to be overcome by the energy
due to the temperature,
and so they're going to be able to break free of each other,
and essentially the helium is going to boil
and you're going to enter into a gaseous state,
the state that most of us are used to seeing helium in.
But that doesn't happen for argon until a good bit warmer.
Still cold by our standards,
and that's because it takes more energy
to overcome the London dispersion forces of argon,
because the argon atoms have larger electron clouds.
So generally speaking, the larger the molecule,
because it has a larger electron cloud,
it'll have higher polarizability
and higher London dispersion forces,
but also the shape of the molecule matters.
The more that the molecules can
come in contact with each other,
the more surface area they have exposed to each other,
the more likely that they can
induce these dipoles in each other.
For example, butane can come in two different forms.
It can come in what's known as n-butane,
which looks like this, so you have four carbons
and 10 hydrogens, so two,
three, four, five, six, seven, eight, nine, 10.
This is known as n-butane,
but another form of butane known as iso
butane would look like this.
So you have three carbons in the main chain,
then you have one carbon that breaks off
of that middle carbon, and then they all have four bonds
and the leftover bonds you could say
are with the hydrogens.
So it would look like this.
This right over here is iso butane.
Iso butane.
Now, if had a sample of a bunch
of n-butane versus a sample of a bunch of iso butane,
which of these do you think will have
a higher boiling point?
Pause this video and think about that.
Well, if you have a bunch of n-butanes next to each other,
imagine other n-butane right over here.
It's going to have more surface area to its neighboring
butanes, because it is a long molecule.
It can expose that surface area to its neighbors.
While the iso butane in some ways
is a little bit more compact.
It has lower surface area.
Doesn't have these big long chains,
and so because you have these longer n-butane molecules,
you're going to have higher London dispersion forces.
They obviously have the same number of atoms in them.
They have the same number of electrons in them,
so they have similar size electron clouds.
They have the same molar mass,
but because of n-butanes' elongated shape,
they are able to get closer to each other
and induce more of these dipoles.
So just by looking at the shape of n-butane versus iso
butane, you say higher London dispersion forces in n-butane,
so it's going to have a higher boiling point.
It's going to require more energy
to overcome the London dispersion forces
and get into a gaseous state.


London dispersion forces | Intermolecular forces and properties | AP Chemistry | Khan Academy

29 分類 收藏
林宜悉 發佈於 2020 年 3 月 28 日
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