Boiling point elevation

The boiling point of a substance is the temperature at which the vapor pressure of the liquid equals the pressure surrounding the liquid and the liquid changes into a vapor.

The boiling point of a liquid varies depending upon the surrounding environmental pressure. A liquid in a partial vacuum has a lower boiling point than when that liquid is at atmospheric pressure. A liquid at high pressure has a higher boiling point than when that liquid is at atmospheric pressure. For a given pressure, different liquids boil at different temperatures. For example, water boils at 100 °C (212 °F) at sea level, but at 93.4 °C (200.1 °F) at 2,000 metres (6,600 ft) altitude.

The normal boiling point (also called the atmospheric boiling point or the atmospheric pressure boiling point) of a liquid is the special case in which the vapor pressure of the liquid equals the defined atmospheric pressure at sea level, 1 atmosphere. At that temperature, the vapor pressure of the liquid becomes sufficient to overcome atmospheric pressure and allow bubbles of vapor to form inside the bulk of the liquid. The standard boiling point has been defined by IUPAC since 1982 as the temperature at which boiling occurs under a pressure of 1 bar.

The heat of vaporization is the energy required to transform a given quantity (a mol, kg, pound, etc.) of a substance from a liquid into a gas at a given pressure (often atmospheric pressure).

Liquids may change to a vapor at temperatures below their boiling points through the process of evaporation. Evaporation is a surface phenomenon in which molecules located near the liquid's edge, not contained by enough liquid pressure on that side, escape into the surroundings as vapor. On the other hand, boiling is a process in which molecules anywhere in the liquid escape, resulting in the formation of vapor bubbles within the liquid.

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 Macroscopic View

When a solute is added to a solvent, the vapor pressure of the solvent (above the resulting solution) is less than the vapor pressure above the pure solvent. The boiling point of a solution, then, will be greater than the boiling point of the pure solvent because the solution (which has a lower vapor pressure) will need to be heated to a higher temperature in order for the vapor pressure to become equal to the external pressure (i.e., the boiling point).

The boiling point of the solvent above a solution changes as the concentration of the solute in the solution changes (but it does not depend on the identity of either the solvent or the solute(s) particles (kind, size or charge) in the solution).

Non-Volatile Solutes

The boiling point of the solvent above a solution will be greater than the boiling point of the pure solvent whether the solution contains a non-volatile solute or a volatile solute. However, for simplicity, only non-volatile solutes will be considered here.

Experimentally, we know that the change in boiling point of the solvent above a solution from that of the pure solvent is directly proportional to the molal concentration of the solute:

T = K_bm

where:

        T is the change in boiling point of the solvent,
        K_b is the molal boiling point elevation constant, and
        m is the molal concentration of the solute in the solution.

Note that the molal boiling point elevation constant, K_b, has a specific value depending on the identity of the solvent.

solventnormal boiling point, ^oCK_b, ^oC m^-1

water100.00.512

acetic acid118.13.07

benzene 80.12.53

chloroform 61.33.63

nitrobenzene 210.95.24

The following graph shows the normal boiling point for water (solvent) as a function of molality in several solutions containing sucrose (a non-volatile solute). Note that the normal boiling point of water increases as the concentration of sucrose increases.

The Microscopic View

The figure below shows a microscopic view of the surface of pure water. Note the interface between liquid water (below) and water vapor (above).

Non-Volatile Solutes

The figures below illustrate how the vapor pressure of water is affected by the addition of the non-volatile solute, NaCl.

Note that:

there are fewer water molecules in the vapor (i.e., lower vapor pressure) above the NaCl solution than in the vapor above pure water, andthe boiling point of the NaCl solution will be greater than the boiling point of pure water.

Pure water - microscopic view.
Normal boiling point = 100.0^oC.1.0 M NaCl solution - microscopic view.
Normal boiling point = 101.0^oC.
Note that the ionic solid, NaCl, produces Na⁺ ions (blue) and Cl^- ions (green) when dissolved in water.