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Course: Chemistry library > Unit 11
Lesson 1: States of matter- States of matter
- States of matter follow-up
- Specific heat and latent heat of fusion and vaporization
- Specific heat, heat of fusion and vaporization example
- Chilling water problem
- Change of state example
- Vapor pressure
- Phase diagrams
- Representing solids, liquids, and gases using particulate models
- Crystalline and amorphous polymers
- Representing alloys using particulate models
- Structure of metals and alloys
- Solids, liquids, and gases
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Vapor pressure
Vapor pressure, volatility, and evaporation. Created by Sal Khan.
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- In nature, why does air exist continuously in gaseous form ? How does it maintain it's high KE, and why do molecules of water which evaporate condense back again when the air can maintain it's KE ?? Sorry if this is a stupid question.(29 votes)
- air is constitutent of many different elements but mostly nitrogen, oxygen, co2 and varying degrees of water vapour and noble gases. all of these have very weak intermolecular forces even though the forces within the actual atoms may be strong, in nitrogen's case, it is diatomic and has a triple bond with another nitrogen so the atomic forces are strong but the forces between different molecules of nitrogen is almost negligle so it takes little energy to break these and leave it in a gaseous form. therefore the little KE within the atmosphere is sufficent, do remember this is also dependant upon the pressure as well as KE (heat energy). Have a Good Day(31 votes)
- So just for clarification vapor pressure can ONLY occur in a closed system?(10 votes)
- The answer to this is trickier than you might think. We measure vapor pressure in closed systems where a pretty good equilibrium can be established. But we also experience vapor pressure in everyday life. It affects the rate at which water evaporates from our skin. If you put a liquid into a sealed container, a portion of it will quickly evaporate until the vapor pressure is attained. When we sweat, we cool off as a result of water evaporating. If we stay relatively still, the water vapor lingers near our skin and slows the rate of evaporation (because of vapor pressure). If we sit in front of a fan, the water vapor moves away quickly and sweat continues to evaporate. Fans don't lower the temperature of the room, but they allow our sweat to evaporate more quickly and make us feel cooler. So vapor pressure can build up in closed systems or those where mixing is slow.(18 votes)
- So, if youre going up a hill, so atmospheric pressure decreases.......will the boiling point change?? If it does, how will it change?(7 votes)
- For example, when you are in the top of mount everest atmospheric pressure is about 26 kPa (on the sea level it's about 101 kPa)and so you don't need 100C to boil water. There the boiling point is 69C.(17 votes)
- What would actually happen if a vacuum was created over the liquid in a closed system? Would the vapor pressure increase?(3 votes)
- 1.) The vapor pressure is a property of the substance and is constant at a given temperature. It increases when temperature increases.
2.) The boiling point of a substance is the temperature at which the vapor pressure of the liquid equals the pressure surrounding the liquid.
-If you create a vacuum over the liquid you decrease the pressure surrounding the liquid. As soon as it falls below the vapor pressure the liquid will start to boil (you do not have to increase the temperature)!
-If you seal the container, the pressure over the liquid will always be the same as the vapor pressure. Thus no matter how much you heat your liquid in order to increase the vapor pressure it will never boil (provided you have a perfect seal and your container can withstand the pressure i.e. it doesn't blow up!). I.E. your boiling point will increase if you seal the container.(10 votes)
- How can I calculate the vapor pressure for methane and ethanol?(2 votes)
- You can use the Antoine Equation . Log (P) = A ' (B / (T + C))
P = vapor pressure (bar), T = temperature (K)
The A, B and C are the Antoine Coefficients and can be found at National Institute of Standards and Technology (NIST) Chemistry Webbook(8 votes)
- How does temperature give and take kinetic energy?(3 votes)
- Energy, such as heat which we measure with a thermometer as temperature, is shown by motion of atoms. As they gain energy they move faster, similarly as you remove heat they move slower. as kinetic energy is the energy of motion, at higher temperatures molecules have on average more motion and energy, while at lower temperatures the molecules on average have less energy and less motion.(4 votes)
- I had an odd question ... That if we need to raise the temperature of the liquid to increase it's vapor pressure , match it with the atmospheric pressure and make it boil , then how does the entire bulk of a liquid kept in open air evaporate ?(3 votes)
- That's because only a few surface molecules have the required amount of kinetic energy to vaporise, we add heat to the liquid to add kinetic energy to the molecules so that an even larger amount of them vaporise. The latent heat of vaporisation is obtained from the surroundings and this makes the process really slow.(2 votes)
- can you please explain someyhing about henry's law?
also, how is teamperature connected to the average kinetic energy?(4 votes)- Henry's Law: The law says that, at equlibrium, the partial pressure (p) of a dissolved substance is proportional it's concentration (c) in the liquid phase; mathematically, that is p = H*c, where H is Henry's constant, which is different for every substance and varies with temperature. Henry's law works quite well as long as the concentration of dissolved substance is low.
Temperature and kinetic energy: For a monatomic ideal gas, the kinetic energy is 12.47 J/(mol,K). See http://en.wikipedia.org/wiki/Kinetic_theory#Temperature_and_kinetic_energy.(1 vote)
- So every time I add a container of a new kind of liquid to a sealed container, the pressure in the container increases? Is the vapor pressure of a liquid only dependent on the pressure that is created by gas particles from exactly that liquid? If so, wouldn't the pressure in the container increase by just adding more containers with different types of liquids?(2 votes)
- Well, it could be any gas. So, a box with 1 mol of H2 gas might have a pressure of 1atm but when you add 1mol of O2 then you can have 2atm. The type of gas does not matter: it is the number of molecules that affect the pressure.
See also: P*V = n*R*T; so that when n number of molecules increases, and V volume is constant, then pressure must increase. Alternatively in your case, when you add more liquid, then V for gas molecules decreases and since n remains constant, then P must increase.(2 votes)
- Why does water boil at 100+ degrees and sea level atm when the vapor pressure of water does not equal the atmospheric pressure?(2 votes)
- Because it does. For water, the vapor pressure reaches the standard sea level atmospheric pressure of 760 mmHg at 100°C.(2 votes)
Video transcript
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.