In an open container the molecules in the gaseous phase will just fly off and an equilibrium would not be reached, as many fewer gaseous molecules would be re-entering the liquid phase.
Also note that at equilibrium the movement of molecules between liquid and gas does not stop, but the number of molecules in the gaseous phase stays the same—there is always movement between phases.
So, at equilibrium there is a certain concentration of molecules in the gaseous phase; the pressure the gas is exerting is the vapor pressure. As for vapor pressure being higher at higher temperatures, when the temperature of a liquid is raised, the added energy in the liquid gives the molecules more energy and they have greater ability to escape the liquid phase and go into the gaseous phase. Let's say you liked to eat turnip greens but didn't like the smell of them cooking.
What you would want to do is cook them quicker, in that case. To cook your greens you put them in a pot of boiling water In an open container water won't get hotter than that. Rather you'll notice a lot of steam coming out of the pot due to evaporation. Adding more heat won't raise the water temperature, and thus, won't cook your greens faster.
If you wanted to cook your turnip greens quicker, you would want the water temperature to be higher. But, there is a solution that will lessen the time you have to smell the greens. You can use vapor pressure to "trick" your turnip greens by using a closed container to cook in—known as a pressure cooker. Pressure cookers have lids that can be secured to the pot which prevents steam from escaping the pot, which raises the pressure of the vapor inside the container.
There is a pressure-release valve on the top of the pot to prevent pressures from getting so high that the pot explodes although there are many instances of the valve malfunctioning with the disastrous effect being a pot that literally explodes. 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 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. For example, diethyl ether is a nonpolar liquid with weak dispersion forces. Water is a polar liquid whose molecules are attracted to one another by relatively strong hydrogen bonding.
Vapor pressure is dependent upon temperature. When the liquid in a closed container is heated, more molecules escape the liquid phase and evaporate. The greater number of vapor molecules strike the container walls more frequently, resulting in an increase in pressure. The Table below shows the temperature dependence of the vapor pressure of three liquids. Notice that the temperature dependence of the vapor pressure is not linear. Use the link below to answer the following questions:. Skip to main content.
States of Matter. Search for:. Vapor Pressure Learning Objectives Define vapor pressure. Describe the relationship between the intermolecular forces in a liquid and its vapor pressure. Describe the relationship between the vapor pressure of a liquid and temperature.
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