  It was known from experiment that the behaviour of gases diverges from the gas laws as pressure is increased and they approach liquefaction. The gas laws were presumed to apply to ‘ideal gases’ as opposed to real gases. The behaviour of real gases approaches that of ideal gases as their pressure is reduced. The kinetic theory had an explanation for this distinction, for at high pressure the assumptions of the kinetic theory, that the volume of molecules is small compared with the total volume of the gas they form part of and that the time spent in collision is small compared to the time between collisions, become increasingly inaccurate. The theory was able to predict various ways in which a real gas will diverge from the ideal gas laws at high pressures (Van der Waals equation) and these were confirmed by experiments on gases approaching liquefaction.