Given a container of $\ce{CO2}$ at a pressure of 100 bar and 295 K, I check the chart below and find:

  1. I'm within the liquid region
  2. I'm before the critical point (304 K)

My question then is, why can I go purchase a tank filled at 100 bar that contains both liquid and gas? What stops the gas from condensing, since according to this chart, the pressure and temperature map to a liquid phase.

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    $\begingroup$ Why is there water vapour in the atmosphere despite liquid water bing the most stable form an normal conditions? $\endgroup$ – matt_black Sep 8 '20 at 20:34
  • $\begingroup$ Using a better phase diagram, it does not correspond to a liquid: chemistry.stackexchange.com/a/128112/79678. $\endgroup$ – Ed V Sep 8 '20 at 20:46
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    $\begingroup$ A CO2 bottle at 20°C does not have 100 bar, but ~57! Or, if it has 100bar, there is zero gas in it, and it is going to explode if you leave it in the sun for a few minutes! $\endgroup$ – Karl Sep 8 '20 at 21:24

The $\ce{CO2}$ exists in two phases when you buy a partially-filled $\ce{CO2}$ cylinder because there is extra volume (which is needed to allow expansion as the container warms). Therefore, the pressure in the container is not 100 bar, but closer to 80 bar at 293K.

If the container had only enough room for the $\ce{CO2}$(l), then, indeed, there could be no gas phase. But if space is allowed above the liquid, then $\ce{CO2}$ evaporates, filling that space, and the $\ce{CO2}$(g) and $\ce{CO2}$(l) reaches equilibrium pressure.

On the molecular level, consider it an ongoing exchange between the molecules of liquid and gas. For the same reason, there is humidity in the air and oceans full of water, a 1 bar at sea level.

  • $\begingroup$ "But if space is allowed above the liquid, then CO2 evaporates, filling that space, and the CO2(g) and CO2(l) reaches equilibrium pressure." -- But after everything settle sand the bottle comes back up to room temperature, what will that equilibrium pressure be? Is it the vapor pressure of liquid CO2 at that temp? $\endgroup$ – dameshgarm Sep 10 '20 at 0:31
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    $\begingroup$ Yes. Vapor is compressible, liquid not much. $\endgroup$ – DrMoishe Pippik Sep 10 '20 at 1:22
  • $\begingroup$ Thanks DrMoishe. Would it then be to correct to say, if I had 3 bottles at room temperature filled with pure CO2 and nothing else (one where 25% of the volume is liquid, another at 50%, and another at 75%), the pressure of all three is the vapor pressure of liquid CO2 at this temperature? $\endgroup$ – dameshgarm Sep 10 '20 at 4:37
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    $\begingroup$ Yes. The pressure would not go above that until it's 100% liquid (and you apply pressure, e.g., from a piston and spring); and would not go below that until it's 0% liquid, i.e., you're drawing gas from the container. $\endgroup$ – DrMoishe Pippik Sep 10 '20 at 21:29

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