Vapor Pressure

Definition of the vapor pressure: The partial pressure of a substance in equilibrium with its condensed phase (liquid or solid). The pressure of the vapor over a liquid (solid) at equilibrium.

We discussed vapor pressure and/or boiling point in CH301 as it related to intermolecular forces (IMF).  Boiling point and vapor pressure are two sides of the same coin.  The boiling point is the temperature at which the vapor pressure equals the total pressure. When comparing vapor pressures we need to be making comparisons at the same temperature. Thus at room temperature, the substance with the lowest boiling point will have the highest vapor pressure (easiest to get into the gas phase). The substance with the highest boiling point will have the lowest vapor pressure.

Vapor pressure is a liquid property related to evaporation. In the liquid (or any substance) the molecules have a distribution of kinetic energies related to the temperature of the system. Because this is a distribution there will always be a few molecules that have enough kinetic energy to over come the attractive potential energy of the other molecules (the intermolecular force), and escape the liquid into the gas phase. In an open container, these molecules will wander off (diffuse) into the room and out into the atmosphere. Eventually all the liquid will evaporate.

In a closed container, the molecules that evaporate will diffuse around in the gas phase, but eventually some of them collide with the liquid. The kinetic energy of the gas molecules has the same distribution as the liquid because they are the same temperature. So now many of the molecules that collide with the liquid will not have sufficient kinetic energy to overcome the IMF and they will "stick" to the liquid. Therefore in a closed container, molecules are both evaporating (turning into gas) and condensing (turning into liquid). The rate of evaporation depends on the temperature and the IMF. The rate of condensation depends on the temperature, the IMF, and the concentration in the gas phase. The condensation rate will initially be zero since there are no molecules in the gas phase. As the evaporation continues the concentration of the molecules in the gas phase (or partial pressure) will increase. This will cause the rate of condensation to increase. When the rate of condensation is equal to the rate of evaporation, the system will no longer change and you'll have a fixed concentration in the gas phase. This concentration can be characterized as a partial pressure. That partial pressure is the "vapor pressure" (VP) of the liquid. The vapor pressure for any pure liquid then depends only on the temperature and the IMF. As temperature increases, VP increases. The stronger the IMF between the molecules, the more energy is required to overcome them and enter the gas phase. Thus the rate of evaporation will be lower and the VP will be lower. Conversely weaker IMF will lead to higher VP.

VP and boiling point (BP) are also related as the boiling point is the temperature at which the VP is equal to 1 atm. Thus the trend in BP with IMF is related to the trend of VP with IMF. High IMF = difficult to "get into" the gas phase. This equates to low VP (few molecules in the gas phase) and high boiling point (high temperature to get the VP = 1 atm).

For example compare three liquids: water, ethanol (C2H5OH), and diethyl ether (C2H5)2O. They have room temperature vapor pressures of 24 Torr, 65 Torr, and 545 Torr respectively. This is the trend we expect from the IMFs. Water is polar and has strong H-­‐bonds, ethanol is also polar but it has weaker H-­‐bonds (only one H that can H-­‐bond vs. the two that water has), and finally diethyl ether is the weakest as it has no H-­‐bonds (no hydrogens bonded to the oxygen). We would expect diethyl ether to have the lowest boiling point (36.6 °C); ethanol to be next highest (78.4 °C); and finally water to exhibit the highest boiling point (100 °C).

Vapor Pressure Movie

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