Deviation from ideal gas behaviour

Deviation from ideal gas behaviour by a real gas can be generally observed at low temperature and high pressure. Only at low pressure and high temperature, do real gases show behavior nearly ideal. Therefore, normally real gases show a large deviation from ideal behavior.

Causes of deviation from ideal gas behaviour

As we have discussed above, the real gases obey the gas laws or the gas equation only when the pressure is low and the temperature is high. But, if the pressure is high or the temperature is low, real gases show remarkable deviation from the ideal behavior.

Then, what may be the reason for deviation from ideal behavior? Ok, let me remind you that gas equation and gas laws were derived on the basis of postulates of the kinetic molecular theory of gases. Failure of these gas laws indicates that some of the postulates of the kinetic theory are seriously in error. These two postulates are:

1. Volume of gas molecules themselves is negligible as compared to the total volume occupied by the gas and
2. There is no force of attraction between the gases molecules

These postulates are valid almost when the pressure is low and the temperature is high. When the pressure is low, the gas molecules are far from each other and the force of attraction between the molecules becomes negligible.

When gas molecules are far from each other, the volume occupied by the gas becomes so large that the volume occupied by individual gas molecules can be neglected.

Positive deviation and negative deviation

Some gases show positive deviation and some show negative deviation. What do these mean? Actually, a gas showing positive deviation indicates that the gas is less compressible than expected from the ideal gas. In positive deviation, the value of Compressibility(Z) is more than one i.e. Z>1.

Similarly, a gas showing a negative deviation indicates that the gas is more compressible than expected from the ideal gas. In negative deviation, the value of Compressibility(Z) is less than one i.e. Z<1.