Q.Differentiate between Real gas and Ideal gas.

An ideal gas is a hypothetical gas that follows PV = nRT exactly; real gases deviate from this, mainly at high pressure and low temperature, because of the finite volume of gas molecules and the intermolecular forces between them (both ignored in the ideal gas model).

Ideal gas: An ideal (or perfect) gas is a hypothetical gas whose molecules are assumed to (i) occupy negligible volume compared to the volume of the container, and (ii) exert no attractive or repulsive forces on one another. Such a gas would obey the ideal gas equation, PV = nRT, exactly at all temperatures and pressures. No real gas is perfectly ideal, but many real gases behave nearly ideally at low pressure and high temperature.

Real gas: A real gas is an actual gas (like O2, CO2, N2, etc.) that shows measurable deviations from ideal gas behaviour, particularly at high pressure and low temperature, where the ideal gas equation no longer accurately predicts its P-V-T behaviour.

Cause of deviation: the two assumptions of the ideal gas model break down for real gases:

1. Finite molecular volume: real gas molecules do occupy a finite (non-negligible) volume. At high pressure, the gas is compressed into a small volume, so the volume actually occupied by the molecules becomes a significant fraction of the total volume, making the gas harder to compress further than an ideal gas would be. …

Real gas molecules have volume and attract each other, unlike ideal gas molecules; these effects vanish at low pressure and high temperature, where real gases behave almost ideally.

The ideal gas equation (PV = nRT) is derived assuming (i) gas molecules have negligible volume compared to the container, and (ii) there are no intermolecular forces of attraction or repulsion between molecules. Real gases deviate from this behaviour because both assumptions fail to some extent in reality:

1. Finite molecular volume: Real gas molecules do occupy a small but finite volume. At high pressure, gas molecules are pushed close together, so their own volume becomes a significant fraction of the total volume, and the actual free volume available for movement is less than the container volume -- causing deviation.

2. Intermolecular forces: Real gas molecules attract one another (van der Waals forces). At low temperature, molecules move more slowly, so these attractive forces are more effective at pulling molecules together, reducing the pressure the gas exerts compared to an ideal gas -- causing deviation. At low temperature and high pressure, this effect can even cause the gas to liquefy.

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