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
No they are not the same. The saturation temperature is the temperature at which the vapour pressure $p(T)$ of the liquid equals the partial pressure of the vapour in the ambient atmosphere. When $T$ is below saturation the vapour in the atmosphere will condense onto the liquid surface. When $T$ is above saturation the liquid will evaporate from the liquid surface until the local atmospheric vapour pressure has increased to that ($p(T)$) required for equilibrium with the liquid.
The boiling point is the temperature at which the vapour pressure of the liquid equals the total pressure of the surrounding atmosphere. The total pressure is (approximately) the sum of the vapour pressure in the atmosphere together with that of the other gaseos components ($O_2$ and $N_2$ etc.). When the liquid vapour pressure is greater that the total atmospheric pressure, the liquid can turn into vapour throughout the body of the liquid (i.e.boil) rather than merely evaporate from the surface.