# Vapor pressure and heat of vaporization relationship advice

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

Determining vapor pressure curves and enthalpy of vaporization using equation describes the relationship between the vapor pressure of a substance, p, and. To know how and why the vapor pressure of a liquid varies with temperature. Nearly all of .. The relationship between pressure, enthalpy of vaporization, and temperature is given by the Is there a scientific basis for this recommendation?. The vapor pressure of pure liquid ${\mathrm{He}}^{3}$ was measured from and the data are represented accurately by the equation The Vapor Pressure, Critical Point, Heat of Vaporization, and Entropy of Liquid He3 . Submit a Manuscript · Publication Rights · Open Access · Tips for Authors · Professional Conduct.

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.

## Vapor pressure

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.

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.

ALEKS - Calculating Vapor Pressure from Boiling Point and Enthalpy of Vaporization

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: 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.

### Heat of vaporization of water and ethanol (video) | Khan Academy

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.

The method should be routinely applicable, reliable, rapid and straightforward to perform. In this article, we will show that thermogravimetric analysis TGA using a reference substance is a suitable technique.

Theoretical background Several methods are available for determining the vapor pressure of substances. Some are based on techniques such as ebulliometry or gas saturation methods, others on static methods using a manometer e. Another group of methods makes use of the so-called effusion technique.

The effusion technique is based on the determination of mass loss. Devices such as the Knudsen cell and techniques derived from this have become of special interest. This is because their theoretical description takes into account effects that depend on a particular experimental setup.

## Determination of Vapor Pressure and the Enthalpy of Vaporization by TGA

These can be due to different geometrical factors or result from the necessity of not having to work in high vacuum in contrast to the original Knudsen method. An effusion cell is shown schematically in Figure 1. As we've already talked about, in the liquid state and frankly, in the solid state as well, the hydrogen bonding is what is keeping these things together, that's what's keeping the water together, flowing next to each other.

This is what's keeping the ethanol together. So if you have less hydrogen-- Let me write this down, less hydrogen bonding, it actually has more hydrogen atoms per molecule, but if you have less hydrogen bonding, it's gonna take less energy to break these things free.

Before I even talk about breaking things free and these molecules turning into vapors entering their gas state, let's just think about how that happens. When we talk about the temperature of a system, we're really just talking about the average kinetic energy. Each molecule, remember they're all bouncing around in all different ways, this one might have, for example, a much higher kinetic energy than this one.

They're all moving in different directions, this one might have a little bit higher, and maybe this one all of a sudden has a really high kinetic energy because it's just been knocked in just the exact right ways and it's enough to overcome both these hydrogen bonds over here and the pressure from the air above it. Remember this isn't happening in a vacuum, you have air up here, air molecules, I'll just draw the generic, you have different types of things, nitrogen, carbon dioxide, etcetera etcetera.

But if I just draw generic air molecules, there's also some pressure from these things bouncing around but this one might have enough, this particular molecule might have enough kinetic energy to overcome the hydrogen bonds and overcome the pressure from the molecules above it to essentially vaporize, to turn into its gas state.

The same thing might be true over here, maybe this is the molecule that has the super high kinetic energy to be able to break free. In that case, it is going to turn into its gaseous state. The hydrogen bonds are gonna break apart, and it's gonna be so far from any of its sibling molecules, I guess you could say, from the other ethanol molecules that it won't be able to form new hydrogen bonds.

Same thing with this one, once it vaporizes, it's out in gaseous state, it's much further from any other water molecules, it's not going to be able to form those hydrogen bonds with them.