Question #f33e6

Boyle's Law [PV=constant] is an approximation to the real behaviour of gases at a fixed temperature. The approximation works sufficiently well for many gases for a suitably small range of conditions, notably reasonably low pressures and high temperatures, well away from their boiling points, but is useless at high densities, high pressures and low temperatures.

Boyle's Law is a good approximation for fluids that are a gas a room temperature (oxygen, argon, helium, nitrogen and many others). Water vapour at around 100°C and one bar is not going to obey Boyle's Law very closely, nor is helium around 4K. Carbon dioxide should obey Boyle's Law around room temperature but not around several tens of degrees below 0°C.

Obviously gases used as refrigerants (methyl chloride, CFCs, butane, ammonia...) are useful for that purpose precisely because they do not obey Boyle's Law (and a lot of other simple laws), a feature which enables refrigerators and heat pumps to be made small. Similarly gases used in cyclic heat power plants (usually water) are specifically chosen because they do not obey Boyle's Law, but instead condense, so that the part of the cycle (the compressor) becomes small and cheap.

There may also be regions of very low pressure where Boyle's Law doesn't work such as at ultra-high vacuums, because of the tendency of gas molecules to stick to the inside surface of the vacuum chamber. However, that is beyond me and you would have to ask an experimental vacuum physicist!

See also the "combined gas law"[ PV=(m/M)RT], van der Waal's Equation, and the kinetic theory of gases. If I recall correctly, carbon dioxide is often used as an example of a situation where van der Waal's Equation is a useful approximation to the observed behaviour, even near the boiling point.