Boiling point and Vapor Pressure

by Mike Daniel/MJC Chem. 142/Fall 2000

We described the process of vaporization or evaporation as when a liquid changes into a gas. For example

H₂O_(L) ® H₂O_(G)

The reverse process is called condensation. The gaseous liquid is sometimes called a vapor. When you heat water and bubbles start rising to the surface the water is boiling. The temperature of the water will remain the same as long as it is boiling. The normal boiling point of a liquid is the constant temperature at 1 atmosphere of pressure at which the liquid changes into a gas. We specify the pressure because the boiling point varies with pressure.

Liquids and some solids can evaporate even when the temperature is less than the boiling point. When you wash your car on a hot summer day, the wet ground eventually dries up due to evaporation of the water. Water in an open container evaporates until no more water remains. However, if you cover the container, a small quantity vaporizes and then evaporation appears to stop. The reason it stops is because the vapor above the liquid can condense. When the speed of vaporization is equal to the speed of condensation, the quantity of liquid will no longer change. The liquid and gas are in equilibrium. We usually use a double arrow to show that a reaction is going in the forward and backward direction, but due to display problems I’m using an equal sign.

H₂O_(L) = H₂O_(G)

Below the boiling point, evaporation takes place at the surface of the liquid. The average kinetic energy of the liquid particles depends on the temperature. Molecules near the surface have a greater than average kinetic energy, and they may have a kinetic energy greater than the intermolecular forces holding them together. These molecules with a kinetic energy greater than the IMF holding them together can escape into the gas phase. This is how a liquid evaporates. As more and more molecules enter the gas phase in our closed container, some of them can bounce back into the liquid. We described equilibrium as when molecules are entering the gas phase just as fast as they enter the liquid phase. It people were leaving and entering a store at the same speed, the number of people in the store would be constant, even though they are coming and going.

The vapor above the liquid exerts a pressure like any gas. The vapor pressure of a liquid is the pressure exerted by the liquid’s vapor at equilibrium. As temperature increases, the average kinetic energy of the molecules increases. As the average kinetic energy of the molecules increases, more molecules are going to have a kinetic energy greater than the IMF holding them together. Thus more molecules can go into the vapor state, and the vapor pressure increases. The following graph shows, at two different temperatures, the number of molecules on the y axis and kinetic energy of molecules on the x axis. This shows how at a higher temperature, more molecules have enough energy to leave a liquid and go into the vapor phase. Thus the vapor pressure increases.

The graph below shows how vapor pressure increases at temperature increases for several different substances. Since different liquids have different intermolecular forces, at the same temperature, liquid have different vapor pressures. The greater the IMF, the more energy needed to leave the liquid, thus the lower the vapor pressure. From this graph, we can see that water have the greatest IMF, followed by ethyl alcohol and diethyl ether. From the structures of the three substances given below, you should be able to rationalize the ranking of IMF for the three substances.