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The 'boiling point' of a substance is the 'maximum temperature' at which a liquid can remain a liquid. Adding a small amount of 'heat energy' (latent heat of vaporization) can convert the liquid into a gas. A pure liquid may change to a gas at temperatures below the boiling point through the process of evaporation. Any change of state from a liquid to a gas at boiling point is considered vaporization. However, evaporation is a surface phenomenon, in which only molecules located near the gas/liquid surface could evaporate. Boiling on the other hand is a bulk process, so at the boiling point molecules anywhere in the liquid may be vaporized, resulting in the formation of vapor bubbles.
A somewhat clearer (and perhaps more useful) definition of boiling point is "the temperature at which the vapor pressure of the liquid equals the atmospheric pressure."

Contents

Saturation temperature and pressure

Intermolecular interactions

Properties of other elements

See also

References

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 change.
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 substance equals the ambient pressure. Thus the boiling point is dependent on the pressure. Usually, boiling points are published with respect to standard 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 ambient pressure up to the critical point, where the gas and liquid properties become identical. The boiling point cannot be increased beyond the critical point. Like wise, the boiling point decreases with decreasing ambient 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 substance at a certain temperature is known, the normal boiling point (under standard pressure) can be calculated by:

Where $T\_B$ is the boiling point under standard pressure,
$R$ is the ideal gas constant,
$P\_0$ is the vapor pressure at a given temperature,
$T\_0$ is that temperature, and
$Delta\; H\_\{vap\}$ is the heat of vaporization of the substance.
'Saturation Pressure', or ''vapor point'', 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.

Intermolecular interactions

In terms of intermolecular interactions, the boiling point represents the point at which the liquid molecules possess enough thermal energy to overcome the various intermolecular attractions binding the molecules into the liquid (eg. dipole-dipole attraction, instantaneous-dipole induced-dipole attractions, and hydrogen bonds). Therefore the boiling point is also an indicator of the strength of these attractive forces.
The boiling point of water is 100 °C (212 °F) at standard pressure. On top of Mount Everest the pressure is about 260 mbar (26 kPa) so the boiling point of water is 69 °C. (156.2 °F).
For purists with a knowledge of thermodynamics, the ''normal boiling point of water'' 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.

Properties of other elements

The element with the lowest boiling point is helium. Both the boiling points of rhenium and tungsten exceed 5000 K at standard pressure. Due to the experimental difficulty of precisely measuring extreme temperatures without bias, there is some discrepancy in the literature as to whether tungsten or rhenium has the higher boiling point.
(Cf. DeVoe, Howard, ''Thermodynamics and Chemistry.'' Prentice-Hall, 2001)

See also

★ List of elements by boiling point

★ Leidenfrost effect

★ flash point

★ boiling delay

★ critical temperature

★ triple point

★ boiling-point elevation

References

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