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Will a real gas occupy all the space available for it? An ideal gas will expand to fill its container but will a real gas? And what are the changes in a real gas when it expands freely against a vacuum?

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That's my answer to the question, which may not be correct.

Real gas molecules are bound by (attractive) Van der Waal's forces (VdW forces). When molecules move apart, part of their kinetic energy is converted to the potential energy (because Van der Waal's force is attractive), if no external source of energy (i.e. heat reservoir, external pressure sources/work done) is present. (i.e. there is no exchange of energy across the boundaries of the container).

That said, when a real gas expands into a vacuum, though there's no work done on the gas, the temperature of the gas should drop because part of the kinetic energy (which is proportional to the temperature) is converted to the potential energy.

Will the gas fill the whole container? Now first ignore surface effects (i.e. boundary walls of the container). We have to understand that although the molecules are "bound" by VdW forces, which tend to draw the molecules together, the latter tendency is balanced by the thermal motion of the gas molecules, which tends to pull the molecules apart. When the temperature is so low such that the thermal motion "dies out", the molecules may settle into a lower energy state (e.g. lattice/solid)

In most cases, real gases should fill the whole container.

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    $\begingroup$ If it doesn't fill the whole container, it's not a gas. $\endgroup$ – Curt F. Apr 3 '15 at 19:13

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