Lower temperature means less energy input, less energy input means that the molecules of water change from liquid to vapor state with less energy. In other words the change from one form to another poses less difficulty to the input energy, hence "easier."

I don't think so.

Very nearly the same amount of *additional* energy is needed to change liquid water to water vapor regardless of the pressure.

212F water @ 0 psig (1 atmosphere) will take an additional 970 btu of heat input to change 1 pound of that water to steam.

250F water @ 15 psig (2 atmospheres) will still require an additional 970 btu of heat input (roughly) to change 1 pound of this water into water vapor.

Pressure changes affect the temperature of where water will phase transistion (boil) significantly; not the energy required for the phase transition.

But in any case, lowering the surrounding pressure lowers the boiling point, the question is: Why does such effect take place?

My guess is because the attractive forces that hold water molecules together as a liquid or a solid depend in much, much greater part on the external pressure exerted by the molecules surrounding the water than any attractive forces present between molecules of the same water when it's a vapor.

I've read many times that attractive forces between liquid or solid molecules of water disappear when enough heat is added to the solid or liquid to turn it into a vapor.

Chemically why that is, I dunno...