However great the pressure , a gas cannot be liquified above its :
A. Boyle temperature
B. Inversion temperature
C. Critical temperature
D. Room temperature

Hint: A phase equilibrium curve's terminal point is referred to as a critical point in thermodynamics. The liquid-vapour critical point, the point on the pressure-temperature curve that marks the circumstances in which a liquid and its vapour can coexist.

Complete step-by-step answer:A real gas begins to act like an ideal gas at a certain pressure range at a temperature, which is known as the Boyle temperature. Boyle temperature of a real gas depends upon its nature. Above the Boyle temperature, real gases show positive deviations from ideal gas behaviour. It has nothing to do with liquefaction.

The crucial temperature below which a non-ideal gas expanding at constant enthalpy will experience a temperature reduction and above which will experience a temperature increase is known as the inversion temperature in thermodynamics and cryogenics. So it is also not connected with liquefaction.
Room temperature too has nothing in connection with liquefaction.
Critical temperature is the minimum temperature above which the vapour form of a gas cannot be liquefied however great the pressure is. The density and all other characteristics of the liquid and vapour are equal at the critical temperature. Therefore, regardless of how high the pressure may be, a gas cannot liquefy above this temperature.
Temperature above which gas cannot be liquefied is known as critical temperature.
Option ‘C’ is correct

Note: Gas is liquefied by cooling it to a level below its boiling point so that it can be transferred and stored as liquid. It is necessary to operate at extremely low temperatures, or "cryogenic" temperatures. A sophisticated system of industrial-scale operations is used to reach these temperatures.