We know that when we have some substance in a liquid state,
it has enough kinetic energy for the molecules to move past
each other, but still not enough energy for the
molecules to completely move away from each other.
So, for example, this is a liquid.
Maybe they're moving in that direction.
These guys are moving a little bit slower in that direction
so there's a bit of this flow going on, but still there are
bonds between them.
They kind of switch between different molecules, but they
want to stay close to each other.
There are these little bonds between them and they want to
stay close.
If you increase the average kinetic energy enough, or
essentially increase the temperature enough and then
overcome the heat of fusion, we know that, all of a sudden,
even these bonds aren't strong enough to even keep them
close, and the molecules separate and they get into a
gaseous phase.
And there they have a lot of kinetic energy, and they're
bouncing around, and they take the shape of their container.
But there's an interesting thing to think about.
Temperature is average kinetic energy.
Which implies, and it's true, that all of the molecules do
not have the same kinetic energy.
Let's say even they did.
Then these guys would bump into this guy, and you could
think of them as billiard balls, and they transfer all
of the momentum to this guy.
Now this guy has a ton of kinetic energy.
These guys have a lot less.
This guy has a ton.
These guys have a lot less.
There's a huge distribution of kinetic energy.
If you look at the surface atoms or the surface
molecules, and I care about the surface molecules because
those are the first ones to vaporize or-- I
shouldn't jump the gun.
They're the ones capable of leaving if they had enough
kinetic energy.
If I were to draw a distribution of the surface
molecules-- let me draw a little graph here.
So in this dimension, I have kinetic energy, and on this
dimension, this is just a relative concentration.
And this is just my best estimate, but it should give
you the idea.
So there's some average kinetic energy at some
temperature, right?
This is the average kinetic energy.
And then the kinetic energy of all the parts, it's going to
be a distribution around that, so maybe it looks something
like this: a bell curve.
You could watch the statistics videos to learn more about the
normal distribution, but I think the normal
distribution-- this is supposed to be a normal, so it
just gets smaller and smaller as you go there.
And so at any given time, although the average is here,
there's some molecules that have a
very low kinetic energy.
They're moving slowly or maybe they have-- well, let's just
say they're moving slowly.
And at any given time, you have some molecules that have
a very high kinetic energy, maybe just because of the
random bumps that it gets from other molecules.
It's accrued a lot of velocity or at least a lot of momentum.
So the question arises, are any of these
molecules fast enough?
Do they have enough kinetic energy to escape?
And so there is some kinetic energy.
I'll draw some threshold here, where if you have more than
that amount of kinetic energy, you actually have enough to
escape if you are surface atom.
Now, there could be a dude down here who has a ton of
kinetic energy.
But in order for him to escape, he'd have to bump
through all these other liquid molecules on the way out, so
it's a very-- in fact, he probably won't escape.
It's the surface atoms that we care about because those are
the ones that are interfacing directly with
the pressure outside.
So let's say this is the gas outside.
It's going to be much less dense.
It doesn't have to be, but let's assume it is.
These are the guys that kind of can escape into the air
above it, if we assume that there's some air above it.
So at any given time, there's some fraction of the particles
or the molecules that can escape.
So you're next question is, hey, well, doesn't that mean
that they will be vaporized or they will turn into gas?
And yes, it does.
So at any given time, you have some
molecules that are escaping.
Those molecules-- what it's called is evaporation.
This isn't a foreign concept to you.
If you leave water outside, it will evaporate, even though
outside, hopefully, in your place, is below the boiling
temperature, or the normal boiling temperature of water.
The normal boiling point is just the boiling point at
atmospheric pressure.
If you just leave water out, over time, it will evaporate.
What happens is some of these molecules that have unusually
high kinetic energy do escape.
They do escape, and if you have your pot or pan outside
or, even better, outside of your house, what happens is
they escape, and then the wind blows.
The wind will blow and then blow these guys away.
And then a few more will escape, the wind blows and
blows them all away.
And a few more escape, and the wind blows and blows
them all the way.
So over time, you'll end up with an empty pan that once
held water.
Now, the question is what happens if you
have a closed system?
Well, we've all done that experiment, either on purpose
or inadvertently, leaving something outside and seeing
that the water will evaporate.
What happens in a closed system where there isn't wind
to blow away?
So let me just draw-- there you go.
Let's say a closed system, and I have-- it doesn't have to be
water, but I have some liquid down here.
And there's some pressure from the air above it.
Let's just say it was at atmospheric pressure.
It doesn't have to be.
So there's some air and the air has some
kinetic energy over here.
So, of course, do the water molecules.
And some of them start to evaporate.
So some of the water molecules that are up here in the
distribution, they have enough energy to escape, so they
start hanging out with the air molecules, right?
Now something interesting happens.
This is the distribution of the molecules
in the liquid state.
Well, there's also a distribution of the kinetic
energies of the molecules in the gaseous state.
Just like different things are bumping into each other and
gaining and losing kinetic energy down here, the same
thing is happening up here.
So maybe this guy has a lot of kinetic energy, but he bumps
into stuff and he loses it.
And then he'll come back down.
So there's some set of molecules.
I'll do it in another set of blue.
These are still the water-- or whatever the fluid we're
talking about-- that come back from the vapor state back into
the liquid state.
And so what happens is, there's always a bit of
evaporation and there's always a bit of condensation because
you always have this
distribution of kinetic energies.
At any given moment in time, out of the vapor above the
liquid, some of the vapor loses its kinetic energy and
then it goes back into the liquid state.
Some of the surface liquid gains kinetic energy by random
bumps and whatever else and goes into the vapor state.
And the vapor state will continue to happen until you
get to some type of equilibrium.
And when you get that equilibrium, we're at some
pressure up here.
So let me see, some pressure.
And the pressure is caused by these vapor particles over
here, and that pressure is called the vapor pressure.
I want to make sure you understand this.
So the vapor pressure is the pressure created, and this is
at a given temperature for a given molecule, right?
Every molecule or every type of substance will have a
different vapor pressure at different temperatures, and
obviously every different type of substance will also have
different vapor pressures.
For a given temperature and a given molecule, it's the
pressure at which you have a pressure created by the vapor
molecules where you have an equilibrium.
Where you have just as many things vaporizing as things
going back into the liquid state.
And we learned before that the more pressure you have, the
harder it is to vaporize even more, right?
We learned in the phase state things that if you are at 100
degrees at ultra-high pressure, and you were dealing
with water, you would still be in the liquid state.
So the vapor creates some pressure and it'll keep
happening, depending on how badly this
liquid wants to evaporate.
But it keeps vaporizing until the point that you have just
as much-- I guess you could kind of view it as density up
here, but I don't want to think-- you have just as many
molecules here converting into this state as molecules here
converting into this state.
So just to get an intuition of what vapor pressure is or how
it goes with different molecules, molecules that
really want to evaporate-- and so why would a molecule want
to evaporate?
It could have high kinetic energy, so this would be at a
high temperature.
It could have low intermolecular forces, right?
It could be molecular.
Obviously, the noble gases have very low molecular
forces, but in general, most hydrocarbons or gasoline or
methane or all of these things, they really want to
evaporate because they have much lower intermolecular
forces than, say, water.
Or they could just be light molecules.
You could look at the physics lectures, but kinetic energy
it's a function of mass and velocity.
So you could have a pretty respectable kinetic energy
because you have a high mass and a low velocity.
So if you have a light mass and the same kinetic energy,
you're more likely to have a higher velocity.
You could watch the kinetic energy videos for that.
But something that wants to evaporate, a lot of its
molecules-- let me do it in a different color.
Something that wants to evaporate really bad, a lot
more of its molecules will have to enter into this vapor
state in order for the equilibrium to be reached.
Let me do it all in the same color.
So the pressure created by its evaporated molecules is going
to be higher for it to get to that equilibrium state, so it
has high vapor pressure.
And on the other side, if you're at a low temperature or
you have strong intermolecular forces or you have a heavy
molecule, then you're going to have a low vapor pressure.
For example, iron has a very low vapor pressure because
it's not vaporizing while-- let me think of something.
Carbon dioxide has a relatively much
higher vapor pressure.
Much more of carbon dioxide is going to evaporate
when you have it.
Well, I really shouldn't use that because you're going
straight from the liquid to the solid state, but I think
you get the idea.
And something that has a high vapor pressure, that wants to
evaporate really bad, we say it has a high volatility.
You've probably heard that word before.
So, for example, gasoline has a higher-- it's more volatile
than water, and that's why it evaporates, and it also has a
higher vapor pressure.
Because if you were to put it in a closed container, more
gasoline at the same temperature and the same
atmospheric pressure, will enter into the vapor state.
And so that vapor state will generate more pressure to
offset the natural inclination of the gasoline to want to
escape than in the case with water.
Now, an interesting thing happens when this vapor
pressure is equal to the atmospheric pressure.
So right now, this is our closed container and you have
the atmosphere here at a certain pressure.
Let's say until now, we've assumed that the atmosphere
was at a higher pressure, for the most part keeping these
molecules contained.
Maybe some atmosphere molecules are coming in here,
and maybe some of the vapor molecules are escaping a bit,
but it's keeping it contained because this is at a higher
pressure out here than this vapor pressure.
And of course the pressure right here, at the surface of
the molecule, is going to be the combination of the partial
pressure due to the few atmospheric molecules that
come in, plus the vapor pressure.
But once that vapor pressure becomes equal to that
atmospheric pressure, so it can press out with the same
amount of force-- you can kind of view it as force per area--
so then the molecules can start to escape.
It can push the atmosphere back.
And so you start having a gap here.
You start having a vacuum.
I don't want to use exactly a vacuum, but since the
molecules escaped, more and more of these molecules can
start going out.
And at that point, you've reached the boiling point of
the substance when the vapor pressure is equal to the
atmospheric pressure.
Just to get a sense of what all of this means, let's look
at the vapor pressure for water.
This is water right here, H2O.
I should do that in black.
And so you see at 760-- so atmospheric pressure, we're in
torr now, but that's just a different-- 760 torr is equal
to 1 atmosphere, so that's about right.
That's about right there, so it's 1 atmosphere.
So at atmospheric pressure, the vapor pressure at 100
degrees Celsius for water-- the vapor is at 100 degrees
Celsius for water.
Or I guess another way to put it, at 100 degrees Celsius,
you have 760 torr of vapor pressure, which is exactly the
atmospheric pressure, or 1 atmosphere, at sea level.
So at 100 degrees, vapor pressure is equal to
atmospheric, or sea level atmospheric.
And so you're going to boil, which we all know is true.
And then at lower temperatures, your vapor
pressure is going to be lower than the
atmospheric pressure, right?
Let's see, here it looks like 300 something.
But then what happens?
If you lowered the atmospheric pressure enough, if you were
to pump air out of the container, or whatever, low
enough, so if you brought the atmospheric pressure down to
this vapor pressure, then again, you will have boiling.
And we saw that in the phase change diagrams, that you can
boil something at a lower temperature if you lower the
atmospheric pressure.
And that's because you're lowering the atmospheric
pressure to the vapor pressure of the substance.
And here's a comparative chart, and this is
interesting.
You see this is kind of an exponential increase with
temperature of vapor pressure.
And that's because, if you think about that distribution
we did before, this is at one kinetic energy.
If you increase the amount of kinetic energy, then your
distribution will look like this.
The temperature has gone up.
And now you have a lot, lot more.
It's not just linear.
You have a lot more particles that can now escape and have
the kinetic energy to evaporate.
And you can see it's this exponential increase as you
increase the temperature.
Now, here's another chart.
You say, hey, where's that exponential increase going?
That's because this is a logarithmic chart.
You can see the scale.
It increases exponentially as opposed to linearly, so it
goes from 0.1 to 10, so equal distances are actually up by a
factor of 10, so that's why you don't see
that logarithmic move.
But these are just for different substances.
Propane, you see at any given-- so let's go at like a
decent temperature.
Let's go 20 degrees Celsius.
At 20 Celsius, propane has the highest vapor pressure.
So this is 1 atmosphere, so propane will actually
evaporate, will actually boil at 20 degrees Celsius.
It will actually completely boil and go
into the gaseous state.
Because its vapor pressure is so much higher than
atmospheric pressure, if we're assuming we're at sea level.
And you could do that for different molecules.
Methyl chloride is the next one.
It's a slightly lower vapor pressure,
but still very volatile.
It would still definitely boil and turn into the gaseous
state at 20 degrees Celsius if we're at sea level because sea
level is right there.
Let's see, at sea level, if you wanted to keep something--
so sea level is this pressure-- if you wanted to
keep let's say, methyl chloride.
If you wanted to keep methyl chloride in the liquid state,
or in equilibrium with the liquid state instead of
boiling, you would have to be at least at
around-- what is this?
Minus 25 degrees Celsius in order for that.
Propane, even at minus 25 degrees is still in the
gaseous state because its vapor
pressure is still higher.
And then, of course, if you have butane, for example.
Butane I think is what they put in lighters, but butane
will be in the liquid state as long as you're at around
roughly 0 degrees.
In a lighter, you might say, oh, it's in a liquid state.
They probably increase the pressure so the pressure in
the lighter is probably something higher.
Maybe it's at 2 atmospheres or something, so that the butane
at room temperature will stay in the liquid state.
Who knows?
I don't know what the pressure is in there.
This is just an interesting chart to look at, that there's
actually a bunch of different vapor pressures.
You can see at atmospheric pressure what's likely to be a
gas or a liquid at different temperatures, and then you
could see at different temperatures, which are the
things that are most volatile and how much do you have to
increase or decrease the pressure to evaporate
something or to boil it.
Anyway, hopefully you found that useful.
Vapor pressure is something that we encounter every day,
and I'll see you in the next video.