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I have seen some phase diagrams that look like this: https://www.learner.org/courses/chemistry/images/text_img/phase_diagram_water.jpg

And also some like this: https://en.wikipedia.org/wiki/Phase_diagram#/media/File:Phase-diag2.svg

My confusion is related to the Critical Point. The first picture doesn't really explain where the gas/supercritical fluid boundary is, and the right angle indicating the supercritical fluid boundary at the critical point seems unusual to me. I have some general questions about behavior beyond the critical temperatures and pressures.

I think my specific question is best summarized by this picture. enter image description here

More specifically, suppose you have an arbitrary compound X with a Liquid/Gas critical temperature of 320 K and critical pressure of 40 Bar which follows the general shape of the above phase diagrams. I have the following questions

1) At 330 K and 10 bar (basically beyond the critical temperature but below the critical pressure) would X be a gas? or would it be a supercritical fluid? Would we need more information to be able to decide?

2) As a more general part of 1), is it possible that a gas exists beyond its critical temperature?

3) At pressures above its critical pressure, a liquid can still exist so long as it's temperature is lower than the critical temperature correct?

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  • $\begingroup$ All that really happens at the critical point is that the liquid and vapor have the same properties. I prefer T-v diagrams myself, then you can see clearly that at the critical point (and above) the densities of the liquid and vapor converge and only one "phase" exists. It is not accurate to call it a liquid or a vapor ... it would be closest to a dense vapor I guess. I don't like your diagram as it seems to insinuate there are phase boundaries above the critical point. Dotted lines would be better to distinguish the supercritical fluid from the vapor or compressed liquid phases. $\endgroup$
    – B. Kelly
    Commented Mar 12, 2019 at 0:45
  • $\begingroup$ I like the second image you linked, that describes how I think of the supercritical region. $\endgroup$
    – B. Kelly
    Commented Mar 12, 2019 at 0:50
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    $\begingroup$ To answer the questions on your diagram, the vapor in that region is still a gas, but is often referred to as superheated. The liquid is still a liquid but is referred to as a compressed liquid. On the phase line, there is just enough pressure to form a liquid... if pressure was decreased some of the liquid would become a vapor.. In the compressed region there is so much pressure that you could decrease the pressure, and it would still stay a liquid... this is a (overly)simple way of thinking about it. $\endgroup$
    – B. Kelly
    Commented Mar 12, 2019 at 0:56
  • $\begingroup$ Often, the term used past the critical point is ‘fluid’. $\endgroup$
    – Jon Custer
    Commented Mar 13, 2019 at 22:25
  • $\begingroup$ Possible duplicate of Why must both the critical temperature and pressure be exceeded to achieve the supercritical phase? $\endgroup$
    – A.K.
    Commented Mar 14, 2019 at 0:26

1 Answer 1

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In order to be in the "critical state", both temperature (x-axis) and pressure (y-axis) must be above the critical point. So answers to your questions become quite obvious:

  1. Very hot but below the critical pressure, it is a gas. You don't need any more information than the temperature and the pressure.
  2. Yes if its pressure is less than its critical pressure
  3. Correct.
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