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If water is introduced in a container maintained at 20 °C in vacuum conditions, a gaseous phase will appear and the pressure will stabilize at the vapour pressure for the given temperature inside the container. So far so good.

Now imagine the experiment is repeated but instead of vacuum conditions, the water is presurized with nitrogen at 1 atm. According to the phase diagram of water, liquid is the stable form of water in these conditions. Yet it is commonly observed that the water molecules with the highest kinetic energy will escape and form a gaseous phase. The partial pressure of gaseous water will be equal to the saturation pressure at this temperature.

Why do we say that this is a liquid/vapour equilibrium in this case, since the liquid phase is not at pressure required for this equilibrium to appear?

Edit: I just realized that my question seems to be answered by Raoult's law.

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When you say “not at the pressure required for this equilibrium to appear”, you make this statement by comparison with water’s phase diagram. However, water’s phase diagram tells you of the behaviour of bulk water, and your system including nitrogen is not bulk water, it’s bulk water plus nitrogen. Thus, comparison of the total pressure of this composite system to the water’s bulk phase diagram is not valid.

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