Today I explained why water boils at different temperatures depending on elevation, and a listener asked if that meant Denver was "less hot" than Miami. The question reminded me how easy it is to conflate pressure and temperature when both affect boiling points.
Boiling point is the temperature at which a liquid's vapor pressure equals the surrounding atmospheric pressure. At sea level, water boils at 100°C because that's when its vapor pressure matches standard atmospheric pressure (101.3 kPa). In Denver, at roughly 1,600 meters elevation, atmospheric pressure drops to about 84 kPa, so water boils at 95°C. The water isn't "less hot"—it's just reaching the vapor-pressure threshold at a lower temperature.
I used a pressure cooker analogy. By sealing the pot, you raise the internal pressure above atmospheric, forcing water to boil at temperatures above 100°C. The reverse happens on a mountain: lower external pressure lets molecules escape the liquid phase more easily. In both cases, the phase transition depends on the balance between molecular kinetic energy (temperature) and the pressure pushing down on the surface.