# Relationship between vapor pressure and density

### Vapor pressure (video) | States of matter | Khan Academy

Accuracy: For many substances, vapor pressures are only poorly Gas density assuming a monatomic gas (unless the substance is given as molecule). unreliable (e.g. large difference between vapor pressure of the solid. The vapor pressure (P°) is the pressure of the vapor of a compound in Normal boiling point: temperature at which the liquids vapour pressure equals 1 atm. A New Relation between Viscosity, Vapor Pressure and Density. Frank Hovorka. J. Am. Chem. Soc., , 55 (12), pp – DOI: /jaa

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?

### Vapor Pressure Calculator [IAP/TU Wien]

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.

## Vapor pressure

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.

### A Relationship between Vapor Pressure and Viscosity of Liquid Foods

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.

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. A substance with a high vapor pressure at normal temperatures is often referred to as volatile.

The pressure exhibited by vapor present above a liquid surface is known as vapor pressure. As the temperature of a liquid increases, the kinetic energy of its molecules also increases. As the kinetic energy of the molecules increases, the number of molecules transitioning into a vapor also increases, thereby increasing the vapor pressure. The vapor pressure of any substance increases non-linearly with temperature according to the Clausius—Clapeyron relation.

The atmospheric pressure boiling point of a liquid also known as the normal boiling point is the temperature at which the vapor pressure equals the ambient atmospheric pressure. With any incremental increase in that temperature, the vapor pressure becomes sufficient to overcome atmospheric pressure and lift the liquid to form vapor bubbles inside the bulk of the substance.

Bubble formation deeper in the liquid requires a higher temperature due to the higher fluid pressure, because fluid pressure increases above the atmospheric pressure as the depth increases. More important at shallow depths is the higher temperature required to start bubble formation.

The surface tension of the bubble wall leads to an overpressure in the very small, initial bubbles.