How does molecular weight affect vapor pressure? | Socratic
Homework Help: Calculating mass of vapor from vapor pressure m at 25 degrees C. The vapor pressures are kPa, kPa, and Pa. The vapour pressure of a solution is always lower than that of the pure lower molecular mass should cause less vapor pressure depression. Raoult's Law relationship between vapor pressure and concentration of a Fractional vapor pressure lowering can be used to calculate molecular mass.
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?
Vapor pressure (video) | States of matter | Khan Academy
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.
How does molecular weight affect vapor pressure?
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. 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?
Raoult's law, a relation between partial vapour pressure and mole fraction
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. In this section, we describe vapor pressure in more detail and explain how to quantitatively determine the vapor pressure of a liquid. As for gases, increasing the temperature increases both the average kinetic energy of the particles in a liquid and the range of kinetic energy of the individual molecules. The fraction of molecules with a kinetic energy greater than this minimum value increases with increasing temperature.
Just as with gases, increasing the temperature shifts the peak to a higher energy and broadens the curve. Some molecules at the surface, however, will have sufficient kinetic energy to escape from the liquid and form a vapor, thus increasing the pressure inside the container. As the number of molecules in the vapor phase increases, the number of collisions between vapor-phase molecules and the surface will also increase.
Eventually, a steady state will be reached in which exactly as many molecules per unit time leave the surface of the liquid vaporize as collide with it condense. At this point, the pressure over the liquid stops increasing and remains constant at a particular value that is characteristic of the liquid at a given temperature. The rate of evaporation depends only on the surface area of the liquid and is essentially constant.
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. Thus diethyl ether ethyl etheracetone, and gasoline are volatile, but mercury, ethylene glycol, and motor oil are nonvolatile.