Phase diagrams | States of matter and intermolecular forces | Chemistry | Khan Academy

Phase diagrams | States of matter and intermolecular forces | Chemistry | Khan Academy

September 21, 2019 56 By Kody Olson


All of the phase changes we’ve
been doing so far have been under constant pressure
conditions, and, in particular, with the problems
that I’ve been doing with water phase changes in the last
couple of videos, it was at atmospheric pressure, at
least at sea level atmospheric pressure, or at 1 atmosphere. So it was done– well,
I’ll explain this diagram in a second. But we all know that in the
universe, pressure isn’t always constant and it
definitely isn’t always constant at 1 atmosphere. 1 atmosphere was defined
as the pressure at sea level on Earth. Obviously, pressure will vary
wildly if you go to smaller planets or larger planets, or
have thicker atmospheres, or if we’re just doing different
types of applications dealing with gases and liquids
and solids. So what I’ve drawn here
is a phase diagram. Let me write that down. And there are many forms of
phase diagrams. This is the most common form that you might
see in your chemistry class or on some standardized
test, but what it captures is the different states of matter
and when they transition according to temperature
and pressure. This is the phase diagram
for water. So just to understand what’s
going on here, is that on this axis, I have pressure. On the x-axis, I have
temperature, and at any given point, this diagram will tell
you whether you’re dealing with a solid, so solid will
be here, a liquid will be here, or a gas. For example, if I told you that
I was at 0 degrees, let’s say 0 degrees is right there,
if I’m at 0 degrees Celsius and 1 atmosphere, where am I? So 0 degrees, 1 atmosphere,
I’m right at that point right there. So I’m at a boundary point
between solids and liquids at 1 atmosphere of pressure,
right? This is when we’re at 1
atmosphere of pressure. So this coincides with our
traditional notion of when ice freezes or when it melts
at 0 degrees. If we made the pressure
higher, what happens? Well, then ice starts melting at
a lower temperature, right? So this is pressure going up,
so pressure going up, let’s say– I don’t know
what this is. This is maybe 10 atmospheres,
ten times Earth’s atmospheric pressure at sea level, then
all of a sudden, the temperature at which solid
turns into liquid– this transition is solid to liquid
–the temperature at which that happens will go down. Likewise, if we lower the
pressure, if we go to Denver and it’s a mile high, pressure
is lower because we have less of the atmosphere above us,
then all of a sudden, the freezing point increases, so
the freezing point will be something above 1 degree. This isn’t drawn completely to
scale, but the idea is your ice would actually freeze a
little bit faster and would freeze at a higher temperature
in Denver than it would at the bottom of the Dead Sea or in
Death Valley at some below sea level point on the planet. Now, this transition is
the transition between anything and gas. And we’re very familiar,
this is 1 atmosphere. And remember, this is water
we’re dealing with. This is the diagram for water,
so at 1 atmosphere, this is kind of the stuff that
we’re used to seeing. Let me draw a line here. So at 1 atmosphere, 0 degrees is
where solid, or ice, turns into liquid water. And then we go up here, so we
keep going at a higher, higher, higher temperature, and
then here, this would be, since we’re at 1 atmosphere,
this is 100 degrees Celsius right there. And that’s the point at 1
atmosphere of pressure where liquid turns into gas,
or water vaporizes, or the liquid boils. All of those are acceptable
ways to think about that. But what happens when we
go to low pressure? Once again, let’s take our
little trip to Denver. So that’s Denver right there. It’s not that drastic. I’m just doing that for
education purposes. Or even better, let’s say Mount
Everest. Mount Everest, very low pressure there. Then our freezing point, we
already said that goes up when you lower the pressure, and your
boiling point goes down, so it’s much easier to boil
something on the top of Mount Everest than it is to boil it at
the bottom or at the lowest point in Death Valley
or the Dead Sea. The intuition behind that is if
I have a liquid, a bunch of molecules in liquid form, and
they’re touching each other, but they have enough kinetic
energy to move past each other, so they’re flowing past
each other, they’re kind of rubbing up against each other,
one of the reasons why they don’t just evaporate, why this
guy doesn’t just jump up there, is that there’s
air above him. There’s air pressure. And air pressure, we’ve
learned about this when we did PV nRT. That’s a bunch of gas molecules,
and the pressure they’re creating is essentially
caused by their temperature and their
kinetic energy. And they sit there, and they
bounce, and they essentially keep these heavier molecules
from going up. They keep them from essentially
separating from each other and turning
into a gas. So the more pressure you have,
the harder it is for these guys to escape. On the other hand, if we’re in
a vacuum, if we’re doing this on the surface of the moon and
there’s none of these guys there, then just a little
slight bump. Even though this guy’s still a
little bit attracted to over here, they’re still attracted
to each other. But just a little bit of bump,
since there’s no pressure up here on the surface of the moon,
might allow this guy to escape and go straight
to a gas. So when you lower the pressure,
it’s just that much easier to go from liquid to gas
or even from solid to gas. And you might say, Sal,
that’s a bizarre concept, solid to gas. It turns out, if you get to low
enough pressures here, I mean, let’s say this is–
Actually, there’s probably not stuff here. This is probably close to
a vacuum right here. You could go from ice– So if
you took ice and you were on the moon and you were at the
right temperature– this is maybe some negative degrees
Celsius temperature; I don’t know what the exact temperature
is –your ice on the moon would go directly
from ice to a gas. Because there’s this huge
vacuum here, so these molecules would say, hey,
there’s all this space to fill and if they just get bumped a
little bit, they’re just going to escape and turn into a gas. You might say, oh, Sal, that’s
a strange phenomenon. It only exists on the moon. And to rebut that comment, I’ve
drawn the phase diagram for carbon dioxide. It’s all around you. You’re exhaling it
as we speak. Your plants in the room are
hopefully inhaling it, but carbon dioxide at 1 atmosphere
has a very different behavior than water. This is carbon dioxide
at 1 atmosphere. Just so you know, this
scale is definitely not drawn to scale. The difference between 1
atmosphere and 5 atmospheres is not the same as between
5 atmospheres and 73. Likewise, this is not
drawn to scale here. This is a much larger
distance than this. If I had to really draw it to
scale, I’d have to stretch this chart out or do a logarithmic chart or something. But anyway, I was talking
about carbon dioxide. So this is carbon dioxide solid,
and this is gas, and this is liquid carbon dioxide. So at 1 atmosphere, let’s say
you live at sea level, like you’re in New Orleans, I guess
that’s a little bit below sea level– that’s where I grew up
–if you were able to get your fridge down to minus 80 degrees
Celsius, the carbon dioxide would actually freeze. And you’re actually not too
unfamiliar with that, or at least you haven’t been if you’ve
gone to some– I don’t know if they still use it for
smoke machines or for visual effects on stage, but
this is dry ice. It’s frozen carbon dioxide. If you’re at sea level
atmospheric pressure, as soon as you get above this minus 78
and 1/2 degrees Celsius, it sublimates to gas. So that process, where you go
straight from a solid to a gas, is sublimation. And that’s why dry ice, when
you see it, you don’t see liquid dry ice or you don’t see
it at standard pressures. I’ve never seen liquid
carbon dioxide. In fact, to get liquid carbon
dioxide, you have to get above 5 atmospheres so you have to get
above five times the sea level pressure on Earth, and
you’re really not going to see that in natural conditions
on Earth. You might see that on Jupiter
or Saturn where you have tremendous pressures because of
the gravity and all of the atmosphere above you. Liquid carbon dioxide, you might
see– I don’t know if Jupiter actually has carbon, but
you’ll probably see it on other huge massive planets
that are gas giants. But on Earth, this process is
just called sublimation. It’s just a neat word. Or it’s sublimating. It’s going straight from solid
to gas and it’s something you’ve seen with dry ice. Now, there’s a couple other
interesting points here and you’re probably already
noticing them. This right here is called the
triple point, because right here at this– Well, in the case
of carbon dioxide, at 5 atmospheres and minus 56 degrees
Celsius, the carbon dioxide is in a state of
equilibrium between the ice, the liquid and the gas. It’s a little bit of
all of the three. And if you just nudge it in one
direction or another by nudging the pressure
or the temperature, it’ll go in that direction. Similarly, water’s triple
point is right here. It’s at a much lower
pressure than we’re used to dealing with. This is 0.611 kilopascals, or
just 611 pascals, which is 5/1000 of an atmosphere. So if you go down to 5/1000 of
an atmosphere and you go a little bit above 0 degrees
Celsius, you’re at the triple point of water. where water can take on any of
these states if you just nudge it in one direction
or another. Now, the other interesting
point on these charts is up here. This is the critical point. Sounds very important. Critical point. And that’s the point at which
if you increase the temperature beyond that or the
pressure beyond that, you’re dealing with a supercritical
fluid. It sounds very exciting. So above here, you have
a supercritical fluid. So very high temperature,
very high pressure. It’s so high temperature that
it wants to be a gas, but you’re putting so much pressure
on it that it wants to be a fluid, so it’s
a little bit of both. And actually, in the case of
water, supercritical water is actually used as a solvent. Because you can imagine, it’s
kind of like liquid water in that things can dissolve
in it, but it’s so high temperature and it can diffuse
into solids that it’s really good at just getting whatever
you want out of whatever you’re trying to clean or
somehow get into or get salt put into the water. So this is supercritical
fluid and it’s a fun thing to think about. But anyway, I just wanted to
expose you to these phase diagrams. Everything I’ve done
so far was at a constant pressure and I changed the
temperature, but you can also read them the other way. If I’m at 100 degrees, and I go
from– Well, let’s say I’m at 110 degrees, where at sea
level is comfortably in the gaseous phase for– So this
is 110 degrees for water. It’s water vapor. But if I were to increase the
pressure and I keep increasing the pressure and maybe I dig a
hole or something or I go into the ocean, then it’s going to
condense into water or it’s going to condense
into a liquid. If I did that experiment here,
when I increase the pressure, I’m going to reverse
sublimate. And I think I wrote down a
word for what that is. Let me see if I wrote
it down someplace. Oh, no, I didn’t. I didn’t write it down. But essentially it’s something
like condense, but the word is escaping me at the second. It’s something on the word of condensing or falling together. Anyway, I forget the word, but
it’ll go straight from a gas to a solid. So these are pretty neat
diagrams. They actually tell a lot about different substances
and then tell you what happens when the pressure or the
temperature changes.