Deviations from Ideal Gas Behavior

Boyle’s law, Charles’ law, and the other gas laws describe the behavior of real gases satisfactorily at ordinary pressure and temperature. But at low temperatures and high pressures, the behavior of any real gas deviate markedly from ideality.

This can be explained thus: at high pressures, gas molecules are relatively close together, and the actual volume of the molecules themselves become significant compared with the free space between them. As is expected, a gas becomes less compressible at high pressures; therefore, doubling the pressure does not reduce the volume by half.

Under extreme high pressures, the space occupied by the molecules becomes a substantial fraction of the space within which the gas is confined, and the gas resists further decrease in volume, a resistance characteristic of liquid (the gas is liquefied). At low temperatures, all gases become more compressible than would be predicted by the gas laws. This phenomenon is due to attractive forces between molecules, which causes them to crowed closer together.

The attractive forces between non polar molecules are called Van der waals forces - after Johannes Van der waals, the Dutch physicist who postulated their existence in 1873. These forces, although important, are effective only over very short distances. One other type of attractive force (hydrogen bonding), which is stronger than Van der waals forces, can exist between molecules of polar substances.

Polar substances deviate very much from ideal behavior because of the permanent polarity of their molecules. The positive end of each polar molecule attracts the negative end of each of its close neighbors with sufficient effectiveness to keep the substance in the liquid state at temperatures far above the boiling points of substances whose only intermolecular attractive forces are the Van der waals type.

This means that even when substances with only the Van der waals forces of attraction are far in the gaseous state, polar substances remain liquid. Thus, HCl, a polar compound, has a boiling point 101^o higher than that of argon, even though its molecular weight is several units less than the molecular weight of argon. Even more striking is the case of H₂O a polar compound, which has a boiling point 346^o higher than that of neon, although its molecular weight is less than the molecular weight of neon. Because the above two factors (high pressures and low temperatures) which cause non ideal behavior in gases acts in opposition, the net deviation at a particular temperature and pressure depends upon which factor is stronger or prevalent. This vary with different gases. Deviation from ideal behavior, except under extreme conditions, is not serious enough to discredit the kinetic theory of gases.