Simple and Fractional Ditillation
Distillation is a commonly used method for purifying liquids and separating mixtures of liquids into their individual components. Familiar examples include the distillation of crude fermentation broths into alcoholic spirits such as gin and vodka, and the fractionation of crude oil into useful products such as gasoline and heating oil. In the organic lab, distillation is used for purifying solvents and liquid reaction products.
To understand distillation, first consider what happens upon heating a liquid. At any temperature, some molecules of a liquid possess enough kinetic energy to escape into the vapor phase
(evaporation) and some of the molecules in the vapor phase return to the liquid (condensation). An equilibrium is set up, with molecules going back and forth between liquid and vapor. At higher temperatures, more molecules possess enough kinetic energy to escape, which results in a greater number of molecules being present in the vapor phase.
If the liquid is placed into a closed container with a pressure gauge attached, one can obtain a quantitative measure of the degree of vaporization. This pressure is defined as the vapor pressure of the compound, and can be measured at different temperatures.
Consider heating cyclohexane, a liquid hydrocarbon, and measuring its vapor pressure at different temperatures. As shown in the following graph of temperature vs vapor pressure, as the temperature of cyclohexane is increased the vapor pressure also increases. This is true for all liquids. At some point, as the temperature is increased, the liquid begins to boil. This happens when the vapor pressure of the liquid equals the applied pressure (for an apparatus that is open to the atmosphere the applied pressure equals atmospheric pressure (1 atm = 760 mm Hg)). For cyclohexane, this occurs at 81° C. The boiling point (BP) of cyclohexane therefore equals 81° C.
The definition of the boiling point of a liquid in an open container then is the
To understand distillation, first consider what happens upon heating a liquid. At any temperature, some molecules of a liquid possess enough kinetic energy to escape into the vapor phase
(evaporation) and some of the molecules in the vapor phase return to the liquid (condensation). An equilibrium is set up, with molecules going back and forth between liquid and vapor. At higher temperatures, more molecules possess enough kinetic energy to escape, which results in a greater number of molecules being present in the vapor phase.
If the liquid is placed into a closed container with a pressure gauge attached, one can obtain a quantitative measure of the degree of vaporization. This pressure is defined as the vapor pressure of the compound, and can be measured at different temperatures.
Consider heating cyclohexane, a liquid hydrocarbon, and measuring its vapor pressure at different temperatures. As shown in the following graph of temperature vs vapor pressure, as the temperature of cyclohexane is increased the vapor pressure also increases. This is true for all liquids. At some point, as the temperature is increased, the liquid begins to boil. This happens when the vapor pressure of the liquid equals the applied pressure (for an apparatus that is open to the atmosphere the applied pressure equals atmospheric pressure (1 atm = 760 mm Hg)). For cyclohexane, this occurs at 81° C. The boiling point (BP) of cyclohexane therefore equals 81° C.
The definition of the boiling point of a liquid in an open container then is the