#pressure

This morning I made tea at my friend's mountain cabin, and the kettle whistled earlier than I expected. I thought my thermometer was broken—it read only 95°C when the water was clearly boiling. That little moment of confusion reminded me how much we take "100°C" for granted.

Most people think water always boils at 100 degrees Celsius. That's the misconception I carried for years too. But boiling point isn't a universal constant—it's the temperature at which a liquid's vapor pressure equals the surrounding atmospheric pressure. At sea level, atmospheric pressure is about 101.3 kPa, which gives us that familiar 100°C. But change the pressure, and you change the boiling point.

Here's where it clicked for me: imagine you're at 3,000 meters elevation, where atmospheric pressure drops to around 70 kPa. Water boils at roughly 90°C there. The water molecules don't need as much energy to escape into vapor because there's less atmospheric pressure pushing down on the surface. It's like trying to open a door—less resistance means less force required. That's why mountaineers have trouble cooking pasta; it never gets hot enough to cook properly.

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.