Boiling Point Elevation

Boiling point elevation is an identical concept to vapor pressure reduction. Since the vapor pressure of the solution is lower than that of the pure solvent, you need to raise the temperature to an even higher point compared to the pure solvent to get the solution to boil. Conceptually the effect is the same as vapor pressure lowering. The solution has a lower free energy than the pure solvent. This means that it is more stable over a larger temperature range. On a phase diagram this results in an expansion of the "liquid" region which results in an increase in the boiling point (and a decrease in the freezing point). This can be seen on the diagram below:

Quantitatively, the change in temperature can be calculated from a number of factors that include the enthalpy of vaporization, the pure boiling point, and the concentration of the solution. These are typically all wrapped up in a single constant resulting in a simple formula.

\[ \Delta T = iK_b \; m\]

K_b is a constant that depends on the solvent and  m is the total solute concentration in molality. K_b is called the boiling point elevation constant or the ebullioscopic constant.   The little "i" in formula is the van't Hoff factor for how many ions an electrolyte (ionic solute) breaks up into.

This effect is generally very small. Value of K_b depend on the solvent but they are typically in the range of 0.5 - 6 °C molal^-1. For example, the K_b for water is only 0.5 °C molal^-1. So even a highly concentrated 1 M NaCl solution has a boiling point increase of only 1 °C. (i = 2 For water 1 M = 1 m. ΔT = (2)(0.5 °C m^-1)(1 m) = 1 °C)

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