If no amount of pressure can liquefy a gas at critical temperature, why does it liquefy at critical pressure?

Question

I have read that at the critical temperature, no amount of pressure will liquefy a gas but according to the definition of critical pressure, a gas at its critical temperature will liquefy at a critical pressure, this seems to be a contradiction because per the definition of critical temperature "no amount of pressure should be able to liquefy the gas". Please I need your help in understanding this better.

Definitions:

Critical Temperature: The critical temperature of a substance is the temperature at and above which the vapor of the substance cannot be liquefied, no matter how much pressure is applied

Critical Pressure: The critical pressure of a substance is the pressure required to liquefy a gas at its critical temperature. Some examples are shown below. Where I got these definitions from

Answer 1

A couple of issues here. First, there is a simpler confusion between "at" and "above." Second there is a conceptual issue with how liquid/gas is defined in the presence of a critical point.

Above the critical temperature, no amount of no amount of pressure will liquefy a gas. At the critical temperature, if the pressure exceeds the critical pressure, the gas will liquefy.

At atmospheric pressure, when we heat water, there is a phase transition when the water reaches 100C. The water needs extra energy to boil - a non-infinitesimal energy (called latent heat of vaporization) is needed to turn the water into the gas just to get it from 100C to 100.00001C (and a gaseous state). Also the volume of the water at constant pressure increases discontinuously when you go from liquid to gas (obviously). There are many kinds of phase transitions, but they all share the common aspect that the properties of a material radically change when you cross a temperature boundary.

Now let's think about what happens to water's phase transition as we increase the pressure. Well first the boiling temperature increases. But also, as we increase the pressure the phase transition becomes less and less dramatic - the latent heat of vaporization becomes smaller and smaller, and the change in volume across the phase transition becomes smaller. When the pressure reaches the critical pressure the phase transition disappears. The latent heat absorbed by the liquid to cross the liquid/gas transition becomes zero, and there is no longer a discontinuous change in volume.

Beyond the critical pressure, there is no phase transition whatsoever. As you increase the temperature, the properties of the liquid are changing continuously (the volume of the liquid increases continuously).

In fact, you can go from something that's clearly liquid water to something that's clearly gas without experiencing any phase transition by simply going around the critical point. Increase the pressure beyond the critical pressure, then increase the temperature beyond the critical temperature. Then decrease the pressure then decrease the temperature. In every step the properties of the fluid will change continuously, and there will be no moment where the temperature stops changing while latent heat is absorbed. But you still get from liquid to gas.

Answer 2

Look at your phase transition graph , its defined up until teh critical point , so for that critical temp youre gonna have a critical pressure and vice versa.