Real gas and Ideal gas

The structure of a molecule has a significant impact on its behaviour. Even if two compounds have the same number of atoms, their behaviour can be substantially different. Ethanol (C2H5OH) is a transparent liquid with a boiling point of around 79 degrees Celsius. Dimethylether (CH3OCH3) has the same amount of carbons, hydrogens, and oxygens as methyl ether (CH3OCH3), but it boils at a lower temperature (25oC). The distinction is due to the quantity of intermolecular interaction (strong H -bonds for ethanol, weak van der Waals force for the ether).

Ideal Gas and Real Gas

Under all temperature and pressure circumstances, an ideal gas obeys the gas laws. To accomplish so, the gas must strictly adhere to the kinetic-molecular theory. There must be no volume in the gas particles, and no attraction forces between them. There can be no such thing as a perfect gas because none of those conditions can be met. A genuine gas is one that does not behave according to the kinetic-molecular theory’s assumptions. Real gases, fortunately, behave very much like ideal gases when exposed to the temperature and pressure conditions found in a laboratory.

So, under what circumstances do gases behave the least optimally? When a gas is compressed under high pressure, the empty space between the particles shrinks, forcing the molecules closer together. The assumption that the volume of the particles themselves is unimportant becomes less valid as the unoccupied space decreases. The decrease in kinetic energy of the particles leads them to slow down when a gas is cooled. The attractive forces between particles are more noticeable when the particles are travelling at slower speeds. Another perspective is that continuing to cool the gas would eventually change it into a liquid, and a liquid is no longer an ideal gas (see liquid nitrogen in the figure below). In conclusion, at low temperatures and high pressures, a real gas deviates the most from an ideal gas. High temperature and low pressure are the best conditions for gases.

Fig.1: Nitrogen gas has converted to a liquid after being cooled to 77K, and it must be stored in a vacuum-insulated container to avoid vaporising too quickly. 

PV/RT displayed against pressure for 1mol of a gas at three distinct temperatures: 200K, 500K, and 1000K. At all temperatures and pressures, an ideal gas would have a value of 1, and the graph would simply be a horizontal line. As may be observed, there are deviations from the ideal gas. As the pressure rises, the attractive forces lead the gas volume to be less than expected, and the PV/RT value falls below 1. The volume of the particles becomes important as the pressure increases, and the PVRT value rises to larger than 1. It’s worth noting that the size of the departures from ideality is largest at 200K and lowest at 1000K.

Fig.2: At high pressures and low temperatures, real gases differ from ideal gases. 

The strength and type of intermolecular attractive forces that exist between the particles also affect the ideality of a gas. Gases with low attractive forces are preferable to those with high attractive forces. Because neon’s atoms are only attracted by modest dispersion forces, while water vapour’s molecules are attracted by comparatively strong hydrogen bonds, neon is more ideal than water vapour at the same temperature and pressure. Helium is a more perfect gas than neon because its dispersion forces are even weaker than those of neon due to its smaller number of electrons.

Conclusion

We conclude that a  real  gas is one that does not behave according to the kinetic-molecular theory’s assumptions.The features of real gases are discussed, as well as their deviations from ideality.The pressure of the gas rises as the volume of the gas decreases. The pressure lowers as the volume of the gas increases.