# Volume occupied by real gases

Molecules of an ideal gas are assumed to have zero volume, the volume available to them for motion is always the same as the volume of the container. In contrast, the molecules of a real gas have small but measurable volumes. At low pressures, the gaseous molecules are relatively far apart, but as the pressure of the gas increases, the intermolecular distances become smaller and smaller . As a result, the volume occupied by the molecules becomes significant compared with the volume of the container. Consequently, the total volume occupied by the gas is greater than the volume predicted by the ideal gas law. Thus at very high pressures, the experimentally measured value of PV/nRT is greater than the value predicted by the ideal gas law.

This is what I read in a website. In the above statements, what is the volume predicted by the ideal gas law? The volume predicted by the gas law, should be the volume of container. Now if volume occupied by real gas is less than the volume of container, how can the volume occupied by it be greater than that predicted by ideal gas law ? Also, compressibility is given by $Z = \frac{PV}{nRT}$. In this equation is P the measured pressure or the ideal pressure?

When you are "containing" a gas in a container then you are by definition restricting its volume. Lets assume you heat the gas. If the volume is confined then the pressure will increase. Similarly, two different gases will have different pressures if confined to the same volume.

If you somehow control the pressure (there is a movable piston) then a real gas will occupy less volume because of its intermolecular interactions.

Also see this: Volume of different gas in a container