# Behaviour of real gases

Behaviour of real gases  :

Compression factor (Z) is the ratio of the actual molar volume of a gas to the molar volume of a perfect gas under the same conditions.

$\text&space;{Z}&space;=$  $\frac{\text&space;{molar&space;volume&space;of&space;the&space;gas}}{\text&space;{molar&space;volume&space;of&space;a&space;perfect&space;gas}}$  $=$  $\frac{\text&space;{PV}_{\text&space;m}}{\text&space;{RT}}$

For a perfect gas Z = 1. So, deviations of Z from 1 are a measure of how far a real gas departs from ideal behaviour.

For an ideal gas, the graph of Z vs P will be a straight line parallel to the pressure axis.

For gases which deviate from ideality, the value of Z deviates from unity. At very low pressures all gases shown have Z = 1 and behave as an ideal gas. At high pressure, all the gases have Z > 1. These are more difficult to compress. At intermediate pressures, most gases have Z < 1. Thus gases show ideal behaviour when the volume occupied is large so that the volume of the molecules can be neglected in comparison to it. In other words, the behaviour of the gas becomes more ideal when pressure is very low. Upto what pressure a gas will follow the ideal gas law, depends upon nature of the gas and its temperature. The temperature at which a real gas obeys ideal gas law over an appreciable range of pressure is called Boyle temperature or Boyle point. Boyle point of a gas depends upon its nature. Above their Boyle point, real gases show positive deviations from ideality and Z values are greater than one. The forces of attraction between the molecules are very feeble. Below Boyle temperature real gases first show a decrease in Z value with increasing pressure, which reaches a minimum value. On further increase in pressure, the value of Z increases continuously. Above explanation shows that at low pressure and high-temperature gases show ideal behaviour. These conditions are different for different gases.