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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?
Because the boiling point reduces with DECREASING ambient pressure.
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?
