Real gases do not always follow the assumptions of the kinetic molecular theory. While the particles of an ideal gas are assumed to occupy no volume and experience no interparticle attractions, the particles of a real gas do have finite volumes and do attract one another. As a result, real gases are often observed to deviate from ideal behavior. Created by Sal Khan.
- [Instructor] 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, And now we're gonna think a little bit about what does it mean to be ideal and what do real gasses, how do they 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, volume of gas, is negligible, negligible, relative to container, to 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 a 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, and 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. And so that's why these are reasonable assumptions, and they allow us to say that PV is equal to nRT, which is a valuable thing, a valuable approximation in most circumstances. But in the real world, we do know that in actuality, that each molecule takes up some volume, and that if you add up all the molecules together, they're of course going to take up some volume, and if there's enough molecules, or if the container is small enough, we know that the volume of the gas relative to 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 have, you might have some, even for fairly inert molecules, you might have 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 character 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 are negligible relative to the container, and 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, and so in this situation, where we can't make these assumptions, we're going to have to modify the ideal gas law.