On a mountaintop, it is observed that water boils at 90°C, not at 100°C as at sea level. Why does this phenomenon occur on the mountaintop?

Boiling point is the point at which the vapour pressure of the liquid and the pressure of the environment around it, are equal.

At this temperature, because there is no difference between the atmospheric pressure around the liquid, and that of the vapour above it, the liquid changes into vapour phase.

On a mountain top, the atmospheric pressure is somewhat lower than at sea level, because there is a lower "weight" of atmosphere pushing down. Therefore the temperature at which the vapour pressure reaches the same value as atmospheric pressure is somewhat lower than at sea level.

This is why trying to make tea on a very high mountain is often not very successful - the water does boil, but the temperature at which it does so is insufficient to get the flavour and extracts out of the tea!

By definition, the boiling point of a liquid, any liquid, is the temperature at which the vapour pressure of the liquid is equal to the ambient pressure, and bubbles of vapour form directly in the liquid. The #"normal boiling point"# is specified when the ambient pressure (and thus the vapour pressure of the liquid) is equal to #"1 atmosphere."#

This definition is a mouthful, but it does underly the principle of vacuum distillation, and distillation under pressure. If I reduce the ambient pressure, I can distil an otherwise involatile liquid in that I have reduced its boiling point.

As to your problem (finally!), on a mountaintop, the ambient pressure is reduced from #"1 atmosphere"#, and thus we heat to a temperature REDUCED from the #"normal boiling point"# such that the liquid vapour pressure is equal to the reduced ambient pressure, whatever this is.

In Denver, Colorado, you are #1.6*km# above sea level, and water boils at #94# #""^@C#. Is this consistent with what I have argued? Is this at a higher elevation than in your problem?