How can CO2 exist in two different phases?

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.

• Why is there water vapour in the atmosphere despite liquid water bing the most stable form an normal conditions? Sep 8 '20 at 20:34
• Using a better phase diagram, it does not correspond to a liquid: chemistry.stackexchange.com/a/128112/79678.
– Ed V
Sep 8 '20 at 20:46
• 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!
– 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.