Vapour pressure and boiling point relationship tips

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

vapour pressure and boiling point relationship tips

describe the processes of evaporation and condensation. describe equilibrium vapor pressure. express the relationship between boiling point, vapor pressure. An explanation of how the saturated vapour pressure of a pure substance arises and how it varies with temperature. temperature, and the relationship between saturated vapour pressure and boiling point. You can look at this in two ways. To understand that the equilibrium vapor pressure of a liquid depends on the To understand that the relationship between pressure, enthalpy of .. cakes ( cake mixes are often sold with separate high-altitude instructions).

Before it can do so, however, a molecule must also be at the surface of the liquid, where it is physically possible for it to leave the liquid surface; that is, only molecules at the surface can undergo evaporation or vaporization The physical process by which atoms or molecules in the liquid phase enter the gas or vapor phase.

To understand the causes of vapor pressure, consider the apparatus shown in Figure When a liquid is introduced into an evacuated chamber part a in Figure 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 soon as some vapor has formed, a fraction of the molecules in the vapor phase will collide with the surface of the liquid and reenter the liquid phase in a process known as condensationThe physical process by which atoms or molecules in the vapor phase enter the liquid phase. 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 rates of evaporation and condensation over time for a system such as this are shown graphically in Figure 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 equilibriumA state in which two opposing processes occur at the same rate, thus producing no net change in the system. 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 pressureThe pressure exerted by a vapor in dynamic equilibrium with its 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 liquidsA liquid with a relatively high vapor pressure. Although the dividing line between volatile and nonvolatile liquids is not clear-cut, as a general guideline, we can say that substances with vapor pressures greater than that of water Table 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. The equilibrium vapor pressure does, however, depend very strongly on the temperature and the intermolecular forces present, as shown for several substances in Figure Some of the surface liquid gains kinetic energy by random bumps and whatever else and goes into the vapor state.

Vapor Pressure - Chemistry LibreTexts

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.

vapour pressure and boiling point relationship tips

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?

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

It could have high kinetic energy, so this would be at a high temperature. It could have low intermolecular forces, right?

Basic Idea:

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.

  • How are vapor pressure and boiling point related?
  • Vapor pressure
  • Chapter 11.4: 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? 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.

vapour pressure and boiling point relationship tips

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.

vapour pressure and boiling point relationship tips

And you can see it's this exponential increase as you increase the temperature. Now, here's another chart.

11.5: Vapor Pressure

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. But these are just for different substances. Propane, you see at any given-- so let's go at like a decent temperature.

Let's go 20 degrees Celsius. At 20 Celsius, propane has the highest vapor pressure. So this is 1 atmosphere, so propane will actually evaporate, will actually boil at 20 degrees Celsius. It will actually completely boil and go into the gaseous state. Because its vapor pressure is so much higher than atmospheric pressure, if we're assuming we're at sea level.

And you could do that for different molecules. Methyl chloride is the next one. It's a slightly lower vapor pressure, but still very volatile. It would still definitely boil and turn into the gaseous state at 20 degrees Celsius if we're at sea level because sea level is right there.

Let's see, at sea level, if you wanted to keep something-- so sea level is this pressure-- if you wanted to keep let's say, methyl chloride. If you wanted to keep methyl chloride in the liquid state, or in equilibrium with the liquid state instead of boiling, you would have to be at least at around-- what is this?