Water boils at a lower temperature at high altitudes because atmospheric pressure decreases as altitude increases, and a liquid boils when its vapour pressure equals the surrounding pressure. Near sea level under standard atmospheric conditions, water boils at about 100 degrees Celsius because the equilibrium vapour pressure of water reaches approximately one atmosphere at that temperature. At higher elevation, there is less air above the liquid, so the surrounding atmospheric pressure is lower.
Water therefore needs to reach a lower temperature before its vapour pressure becomes equal to the reduced external pressure. The change is not caused primarily by the water becoming intrinsically different at altitude. The intermolecular forces and molecular identity of water remain the same, but the external thermodynamic condition for stable vapour bubble formation changes. This is why a pot of water can boil vigorously at a temperature substantially below 100 degrees Celsius in mountainous regions.
Atmospheric pressure decreases with altitude because the pressure at a given level is produced by the weight of the atmosphere above that level. As elevation increases, the amount of air above a location becomes smaller and atmospheric density generally decreases, so the local pressure falls. The exact pressure depends on altitude, temperature, weather, humidity and atmospheric conditions, so there is no single boiling temperature associated with every altitude.
A rough standard atmosphere gives pressures of about 101.3 kilopascals at sea level, about 84 kilopascals around 1.5 kilometres, about 70 kilopascals around 3 kilometres and about 54 kilopascals around 5 kilometres, though real conditions vary. Since boiling depends on pressure rather than altitude directly, the correct physical explanation is that altitude changes the ambient pressure and the changed pressure alters the temperature at which water's vapour pressure matches the surroundings.
The molecular explanation follows from vapour pressure. Even below its boiling point, some water molecules escape from the liquid surface and enter the gas phase. At a given temperature, the liquid has an equilibrium vapour pressure determined by temperature and molecular interactions. As temperature rises, the fraction of molecules energetic enough to enter the vapour phase increases and the equilibrium vapour pressure rises.
A vapour bubble within the liquid can persist and grow when the vapour pressure inside it is sufficient to balance the surrounding pressure. At lower atmospheric pressure, that condition is achieved at a lower temperature. For example, at an atmospheric pressure of roughly 70 kilopascals, corresponding approximately to a high mountain environment, water boils at a temperature well below 100 degrees Celsius. The exact value must be obtained from water vapour pressure data or an appropriate thermodynamic equation rather than inferred solely from altitude.
The practical consequence is that boiling water at high altitude is cooler than boiling water at sea level. This can affect cooking because many cooking processes depend on temperature as well as time. Foods cooked by boiling or simmering may take longer to reach the desired degree of doneness because the surrounding water cannot become as hot while it remains at its boiling point under the reduced pressure.
Pressure cookers reverse this effect by increasing the pressure above the liquid, which raises the boiling temperature. A pressure cooker can therefore provide cooking temperatures above the normal atmospheric boiling point of water. In laboratory and industrial systems, vacuum chambers can reduce pressure enough to cause water and other liquids to boil at temperatures far below their ordinary boiling points. This principle is also used in vacuum distillation, freeze drying and processes designed to remove water while limiting thermal damage.
The lower boiling temperature at altitude does not mean that water contains less thermal energy simply because it is boiling. Temperature measures an aspect of the thermal state, whereas total internal energy depends on the quantity of substance and its microscopic degrees of freedom. A given mass of boiling water at high altitude has a lower temperature than boiling water at sea level, but it still requires substantial energy to vaporise because the latent heat of vaporisation remains significant.
The latent heat changes with temperature and pressure, so it is not exactly constant across different conditions. The energy needed to heat water from a particular starting temperature to its altitude dependent boiling point may also be lower at higher altitude, but the energy required for the subsequent phase transition remains large. This distinction explains why simply reaching the boiling point does not mean that vaporisation requires little energy.
Weather conditions can cause the boiling point to vary even at the same altitude. Atmospheric pressure changes with the passage of high and low pressure systems, and therefore the boiling temperature of water can change slightly from day to day. Temperature and humidity also affect atmospheric density and pressure profiles, although local pressure is the direct quantity relevant to boiling.
The boiling point can consequently be used as a crude indicator of pressure, and specialised instruments called hypsometers have historically used the known boiling temperature of water to estimate altitude or atmospheric pressure. Modern altitude measurement generally relies on barometric or satellite based methods, but the underlying relationship remains a useful illustration of thermodynamics. A thermometer measuring the water itself provides a direct way to determine the actual boiling temperature under current conditions, while a pressure measurement provides a more general description of the environment.
The phrase "water boils at 100 degrees Celsius" is therefore only conditionally true. More precisely, pure water has a normal boiling point of 100 degrees Celsius at a pressure of one standard atmosphere. At lower pressures it boils below 100 degrees Celsius, and at higher pressures it boils above 100 degrees Celsius. The relationship can be described quantitatively by the Clausius Clapeyron equation over appropriate ranges or by accurate steam tables and equations of state.
The vapour pressure of water increases strongly with temperature, so relatively modest changes in pressure can produce measurable changes in boiling temperature. At sufficiently low pressure, liquid water can boil at temperatures near or below ordinary room temperature. The triple point provides a limiting condition for the existence of liquid water in equilibrium, occurring at approximately 0.01 degrees Celsius and 611.657 pascals. Below the triple point pressure, stable liquid water cannot exist in equilibrium and ice can transition directly to vapour through sublimation.
The high altitude boiling effect illustrates a general thermodynamic principle that phase transitions depend on the surrounding conditions as well as the substance itself. The equilibrium between liquid water and water vapour is governed by Gibbs free energy, and the phase boundary shifts with pressure according to thermodynamic relationships. At the boiling point, the chemical potential of the liquid and vapour phases is equal.
Reducing the pressure changes the point at which that equality occurs and therefore lowers the boiling temperature. The phenomenon is important beyond cooking because pressure dependent boiling affects aircraft systems, mountain environments, meteorology, industrial steam generation, vacuum engineering, food processing and chemical separation. It also illustrates why temperature alone cannot fully specify the state of matter. A statement such as "water is boiling" is incomplete thermodynamically unless the pressure and relevant composition are known, because the same substance can boil at very different temperatures under different external pressures.
