boiling point

boiling point

boiling point, temperature at which a substance changes its state from liquid to gas. A stricter definition of boiling point is the temperature at which the liquid and vapor (gas) phases of a substance can exist in equilibrium. When heat is applied to a liquid, the temperature of the liquid rises until the vapor pressure of the liquid equals the pressure of the surrounding gases. At this point there is no further rise in temperature, and the additional heat energy supplied is absorbed as latent heat of vaporization to transform the liquid into gas. This transformation occurs not only at the surface of the liquid (as in the case of evaporation) but also throughout the volume of the liquid, where bubbles of gas are formed. The boiling point of a liquid is lowered if the pressure of the surrounding gases is decreased. For example, water will boil at a lower temperature at the top of a mountain, where the atmospheric pressure on the water is less, than it will at sea level, where the pressure is greater. In the laboratory, liquids can be made to boil at temperatures far below their normal boiling points by heating them in vacuum flasks under greatly reduced pressure. On the other hand, if the pressure is increased, the boiling point is raised. For this reason, it is customary when the boiling point of a substance is given to include the pressure at which it is observed, if that pressure is other than standard, i.e., 760 mm of mercury or 1 atmosphere (see STP). The boiling point of a solution is always higher than that of the pure solvent; this boiling-point elevation is one of the colligative properties common to all solutions.
Boiling-point elevation describes the phenomenon that the boiling point of a liquid (a solvent) will be higher when another compound is added, meaning that a solution has a higher boiling point than a pure solvent. This happens whenever a non-volatile solute, such as a salt, is added to a pure solvent, such as water.


The boiling point elevation is a colligative property, which means that it is dependent on the presence of dissolved particles and their number, but not their identity. It is an effect of the dilution of the solvent in the presence of a solute. It is a phenomenon that happens for all solutes in all solutions, even in ideal solutions, and does not depend on any specific solute-solvent interactions. The boiling point elevation happens both when the solute is an electrolyte, such as various salts, and a nonelectrolyte. In thermodynamic terms, the origin of the boiling point elevation is entropic and can be explained in terms of the vapor pressure or chemical potential of the solvent. In both cases, the explanation depends on the fact that many solutes are only present in the liquid phase and do not enter into the gas phase (except at extremely high temperatures).

Put in vapor pressure terms, a liquid boils at the temperature when its vapor pressure equals the surrounding pressure. For the solvent, the presence of the solute decreases its vapor pressure by dilution. A non-volatile solute has a vapor pressure of zero, so the vapor pressure of the solution is the same as the vapor pressure of the solvent. Thus, a higher temperature is needed for the vapor pressure to reach the surrounding pressure, and the boiling point is elevated.

Put in chemical potential terms, at the boiling point, the liquid phase and the gas (or vapor) phase have the same chemical potential (or vapor pressure) meaning that they are energetically equivalent. The chemical potential is dependent on the temperature, and at other temperatures either the liquid or the gas phase has a lower chemical potential and is more energetically favourable than the other phase. This means that when a non-volatile solute is added, the chemical potential of the solvent in the liquid phase is decreased by dilution, but the chemical potential of the solvent in the gas phase is not affected. This means in turn that the equilibrium between the liquid and gas phase is established at another temperature for a solution than a pure liquid, i.e., the boiling point is elevated.

The phenomenon of freezing-point depression is analgous to boiling point elevation. However, the magnitude of the freezing point depression is larger than the boiling point elevation for the same solvent and the same concentration of a solute. Because of these two phenomena, the liquid range of a solvent is increased in the presence of a solute.


The extent of boiling-point elevation can be calculated by applying Clausius-Clapeyron relation and Raoult's law together with the assumption of the non-volatility of the solute. The result is that in dilute ideal solutions, the extent of boiling-point elevation is directly proportional to the molal concentration of the solution according to the equation:

ΔTb = Kb · mB


  • ΔTb, the boiling point elevation, is defined as Tb (solution) - Tb (pure solvent).
  • Kb, the ebullioscopic constant, which is dependent on the properties of the solvent. It can be calculated as Kb = RTb2M/ΔHv, where R is the gas constant, and Tb is the boiling temperature of the pure solvent, M is the molar mass of the solvent, and ΔHv is the heat of vaporization per mole of the solvent.
  • mB is the molality of the solution, calculated by taking dissociation into account since the boiling point elevation is a colligative property, dependent on the number of particles in solution. This is most easily done by using the van 't Hoff factor i as mB = msolute · i. The factor i accounts for the number of individual particles (typically ions) formed by a compound in solution. Examples:
    • i = 1 for sugar in water
    • i = 2 for sodium chloride in water, due to the full dissociation of NaCl into Na+ and Cl-
    • i = 3 for calcium chloride in water, due to dissociation of CaCl2 into Ca2+ and 2Cl-

At high concentrations, the above formula is less precise due to nonideality of the solution. If the solute is also volatile, one of the key the assumptions used in deriving the formula is not true, since it derived for solutions of non-volatile solutes in a volatile solvent. In the case of volatile solutes it is more relevant to talk of a mixture of volatile compounds and the effect of the solute on the boiling point must be determined from the phase diagram of the mixture. In such cases, the mixture can sometimes have a boiling point that is lower than either of the pure components; a mixture with a minimum boiling point is a type of azeotrope.

Ebullioscopic constants

Values of the ebullioscopic constants KB for selected solvents:

Compound Boiling point in °C Ebullioscopic constant KB in units of °(C · kg) / mol or C/molal
Acetic acid 118.1 3.07
Benzene 80.1 2.53
Carbon disulfide 46.2 2.37
Carbon tetrachloride 76.8 4.95
Naphthalene 217.9 5.8
Phenol 181.75 3.04
Water 100 0.512


Together with the formula above, the boiling-point elevation can in principle be used to measure the degree of dissociation or the molar mass of the solute. This kind of measurement is called ebullioscopy (Greek "boiling-viewing"). However, since superheating is difficult to avoid, precise ΔTb measurements are difficult to carry out, which was partly overcome by the invention of the Beckmann thermometer. Furthermore, the cryoscopic constant that determine freezing-point depression is larger than the ebullioscopic constant, and since the freezing point is often easier to measure with precision, it is more common to use cryoscopy.

A common mis-attribution of the use of boiling-point elevation is adding salt when cooking foods to elevate the temperature of the water before it boils. However, the temperature increase caused by the amounts of salt added when cooking is generally not enough to raise the temperature by a single degree, as a comparison, seawater has a boiling point of 100.6°C. The salt is added simply to season the food and prevent pasta from sticking.

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