Introduction to real gases | Intermolecular forces and properties | AP Chemistry | Khan Academy

In several other videos, we have talked about the ideal gas law, which tells us that pressure times volume is going to be equal to the number of moles times the ideal gas constant times the temperature measured in Kelvin. Now, in all of our studies of the ideal gas law, we assumed that the gases that we were dealing with were ideal.

Now we're going to think a little bit about what it means to be ideal and how real gases vary from actual ideal gases. Well, in order for us to assume that a gas is ideal, we assume that its volume, the volume that the gas takes up, is negligible relative to the container.

The other thing we assume is that the molecules of the gas don't interact with each other; molecules don't interact. Now, in the real world, we know that all molecules take up some volume, but it could be a reasonable assumption if we're talking about a really huge container and we don't have that high density of molecules in it. It's a reasonable assumption that the volume of the gas itself, that the molecules themselves, are small in volume collectively relative to the container.

It's reasonable in many circumstances to assume that the molecules don't interact. Maybe they don't have strong intermolecular forces. Once again, because they're taking up a small portion of the volume, they might even not get close to each other too often. That's why these are reasonable assumptions, and they allow us to say that PV is equal to nRT, which is a valuable approximation in most circumstances.

But in the real world, we do know that in actuality, the volume that each molecule takes up is some volume, and if you add up all the molecules together, they're, of course, going to take up some volume. If there's enough molecules or if the container is small enough, we know that the volume of the gas relative to the container won't be negligible.

We also know that molecules will interact with each other in some way, shape, or form. Two molecules can't occupy the same space at the same time, so you definitely have some repulsive forces, and you might have, even for fairly inert molecules, some temporary dipoles that get formed—some temporary attraction or some temporary repulsion.

So, if you're dealing with a situation where things are less ideal, and I'm going to make a characterization of it where the molecules are taking up a significant volume relative to the container, you can't say that the volume of the molecules is negligible relative to the container.

We assume that they are interacting with each other. They’re definitely going to repulse each other; they can't occupy the same space at the same time. But they might attract each other at some points or repulse each other at other points. In this situation, where we can't make these assumptions, we're going to have to modify the ideal gas law.