Saturation Temperature and Pressure
A saturated liquid contains as much thermal energy as it can without boiling (or conversely a saturated vapor contains as little thermal energy as it can without condensing).
Saturation temperature means boiling point. The saturation temperature is the temperature for a corresponding saturation pressure at which a liquid boils into its vapor phase. The liquid can be said to be saturated with thermal energy. Any addition of thermal energy results in a phase transition.
If the pressure in a system remains constant (isobaric), a vapor at saturation temperature will begin to condense into its liquid phase as thermal energy (heat) is removed. Similarly, a liquid at saturation temperature and pressure will boil into its vapor phase as additional thermal energy is applied.
The boiling point corresponds to the temperature at which the vapor pressure of the liquid equals the surrounding environmental pressure. Thus, the boiling point is dependent on the pressure. Usually, boiling points are published with respect to atmospheric pressure (101.325 kilopascals or 1 atm). At higher elevations, where the atmospheric pressure is much lower, the boiling point is also lower. The boiling point increases with increased pressure up to the critical point, where the gas and liquid properties become identical. The boiling point cannot be increased beyond the critical point. Likewise, the boiling point decreases with decreasing pressure until the triple point is reached. The boiling point cannot be reduced below the triple point.
If the heat of vaporization and the vapor pressure of a liquid at a certain temperature is known, the normal boiling point can be calculated by using the Clausius-Clapeyron equation thus:
| where: | |
| = the normal boiling point, K | |
| = the ideal gas constant, 8.314 J · K−1 · mol−1 | |
| = is the vapor pressure at a given temperature, atm | |
| = the heat of vaporization of the liquid, J/mol | |
| = the given temperature, K | |
| = the natural logarithm to the base e |
Saturation pressure is the pressure for a corresponding saturation temperature at which a liquid boils into its vapor phase. Saturation pressure and saturation temperature have a direct relationship: as saturation pressure is increased so is saturation temperature.
If the temperature in a system remains constant (an isothermal system), vapor at saturation pressure and temperature will begin to condense into its liquid phase as the system pressure is increased. Similarly, a liquid at saturation pressure and temperature will tend to flash into its vapor phase as system pressure is decreased.
The boiling point of water is 100 °C (212 °F) at standard pressure. On top of Mount Everest, at 8,848 m (29,029 ft) elevation, the pressure is about 252 Torr (33.597 kPa) and the boiling point of water is 71 °C (159.8 °F). The boiling point decreases 1 °C every 285 m of elevation, or 1 °F every 500 ft.
There are two conventions regarding the standard boiling point of water: The normal boiling point is 99.97 degrees Celsius at a pressure of 1 atm (i.e., 101.325 kPa). Until 1982 this was also the standard boiling point of water, but the IUPAC now recommends a standard pressure of 1 bar (100 kPa). At this slightly reduced pressure, the standard boiling point of water is 99.61 degrees Celsius.
Read more about this topic: Boiling Point
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