# Vapor pressure and mass relationship

### Vapour pressure of water - Wikipedia

Homework Help: Calculating mass of vapor from vapor pressure What mass of each substance will be found in the air if there is no ventilation. Download scientific diagram | Relationship between vapor pressure and aromatic hydrocarbons in wastewater followed by gas chromatography mass. The vapor pressure of a liquid is the equilibrium pressure of a vapor above its liquid (or solid); that is, the pressure of the vapor resulting from evaporation of a.

The rate of condensation depends on the number of molecules in the vapor phase and increases steadily until it equals the rate of evaporation.

Equilibrium Vapor Pressure Two opposing processes such as evaporation and condensation that occur at the same rate and thus produce no net change in a system, constitute a dynamic equilibrium. In the case of a liquid enclosed in a chamber, the molecules continuously evaporate and condense, but the amounts of liquid and vapor do not change with time. The pressure exerted by a vapor in dynamic equilibrium with a liquid is the equilibrium vapor pressure of the liquid.

If a liquid is in an open container, however, most of the molecules that escape into the vapor phase will not collide with the surface of the liquid and return to the liquid phase. Instead, they will diffuse through the gas phase away from the container, and an equilibrium will never be established. Volatile liquids have relatively high vapor pressures and tend to evaporate readily; nonvolatile liquids have low vapor pressures and evaporate more slowly.

### How does molecular weight affect vapor pressure? | Socratic

Thus diethyl ether ethyl etheracetone, and gasoline are volatile, but mercury, ethylene glycol, and motor oil are nonvolatile. The equilibrium vapor pressure of a substance at a particular temperature is a characteristic of the material, like its molecular mass, melting point, and boiling point Table It does not depend on the amount of liquid as long as at least a tiny amount of liquid is present in equilibrium with the vapor. Molecules that can hydrogen bond, such as ethylene glycol, have a much lower equilibrium vapor pressure than those that cannot, such as octane.

The nonlinear increase in vapor pressure with increasing temperature is much steeper than the increase in pressure expected for an ideal gas over the corresponding temperature range. The temperature dependence is so strong because the vapor pressure depends on the fraction of molecules that have a kinetic energy greater than that needed to escape from the liquid, and this fraction increases exponentially with temperature.

As a result, sealed containers of volatile liquids are potential bombs if subjected to large increases in temperature. Similarly, the small cans 1—5 gallons used to transport gasoline are required by law to have a pop-off pressure release. Volatile substances have low boiling points and relatively weak intermolecular interactions; nonvolatile substances have high boiling points and relatively strong intermolecular interactions.

Vapor Pressure of Mercury The experimentally measured vapor pressures of liquid Hg at four temperatures are listed in the following table: 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.

Vapor Pressure Basic Introduction, Normal Boiling Point, & Clausius Clapeyron Equation - Chemistry

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 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.

## Vapour pressure of water

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 so atmospheric pressure, we're in torr now, but that's just a different-- 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 degrees Celsius for water-- the vapor is at degrees Celsius for water. Or I guess another way to put it, at degrees Celsius, you have torr of vapor pressure, which is exactly the atmospheric pressure, or 1 atmosphere, at sea level.

So at 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 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.