I'm confused as to what exactly is vapor pressure and saturation pressure. From what I understand, vapor pressure is just the equilibrium pressure of a vapor above a liquid at some temperature. Is this not also the definition of the saturation pressure (ie. an equilibrium pressure at some temperature)? Does a system always tend towards its saturation pressure? Or is it that there is one saturation pressure at some specific temperature, and that upon increasing the temperature, the vapor pressure will eventually reach the saturation pressure? Is the saturation pressure dependent on overall pressure?
[Physics] How exactly does vapor pressure relate to saturation pressure
thermodynamics
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Because liquids are in general much less compressible than gases, meaning that a small reduction in volume requires a big increase in pressure.
But if I choose to apply the same reasoning the other way, given that the saturation pressure for T = 200 ºC is 15,54 bar, I do not understand how this would tell me that the water is vapor.
At T = 200 ºC, your pressure of 5 bar is less than the saturation pressure of 15,54 bar. This means that the water is more compressed; the molecules are closer together (i.e., closer to being a liquid).
And I say still liquid because a lower pressure somehow translates as less energy, and less energy means to me that it still remains liquid.
Adiabatic expansion (lowering pressure without adding heat) will decrease energy, like you've said. This transition, however, is a process; you can't apply this intuition to 2 unrelated states without considering a process between those states.
If you adiabatically expand water (200 ºC, 15,54 bar) to 5 bar, its temperature will drop below 200 ºC. The final state will therefore not be steam.
You can see, therefore, that to go from water (200 ºC, 15,54 bar) to steam at the same temperature (200 ºC, 5 bar), we need to expand while adding heat. You need to add energy to make up for the energy loss of expansion, so that we now have an isothermal expansion. By lowering the pressure, however, we reduce the boiling point of water and a phase change occurs.
You can visualize this better on the $P-v-T$ surface for water below:
Check out the constant temperature line starting from the saturated state to the some lower pressure vapor state (this is our isothermal expansion). You can see that at constant temperatures, we can have steam at lower pressures. This is because it is easier to evaporate water at lower pressure, which leads to your next confusion:
It must increase it's energy (just like with temperature) before going for the phase-change.
Lower pressure actually makes it easier for liquids to evaporate. Can you visualize how molecules would more readily break out of the liquid phase into a vapor under low pressures? There is simply less pressure holding the molecules in place. This is why it's easier to boil water at high altitudes (lower pressure).
Best Answer
You can have vapour when there is no liquid present and that vapour would exert a vapour pressure.
If however you have liquid and vapour present in dynamic equilibrium with one another then the pressure exerted by the vapour is the saturated vapour pressure.
So start off with a container with only vapour in it.
The vapour exerts a vapour pressure.
Now do something eg add liquid, cool the vapour, reduce the volume of the container, etc, so that there is also liquid in the container.
Then the pressure exerted by the vapour is the saturated vapour pressure.