Both vapor pressure and boiling point are affected by the strength of interparticle (intermolecular) forces of attraction. Vapor pressure is defined as the partial pressure of a gas in equilibrium with its liquid at a constant temperature. If we have an open container of liquid, it will eventually evaporate–the gas molecules will just go into the atomosphere. If we have a closed container of liquid, the gas molecules are not able to escape. If the closed container is attached to a manometer, the gas molecules are contained above the liquid and the vapor pressure can be easily determined once equilibrium is achieved.

Equilibrium is dynamic. In this case, at equilibrium, the rate of evaporation is equal to the rate of condensation.

Let’s say we have 2 open beakers — one with water and the other with the same amount of acetone. If the beakers are left for several hours, the acetone will have evaporated while the water is still in the beaker. It will take much longer for the water to evaporate. Notice, water has a boiling point of 100^oC while acetone has a boiling point of 56^oC. The vapor pressure of water is lower than that of acetone. Substances with high vapor pressures are volatile. Acetone is considered to be volatile whereas water is nonvolatile at room temperature.

Vapor pressure and boiling points are both dependent on the strength of intermolecular forces. Water molecules experience intermolecular hydrogen bonding while acetone molecules only experience London dispersion and dipole-diple forces. Because water has stronger intermolecular forces of attractions, it has a higher boiling point and a lower vapor pressure than acetone.

The stronger the intermolecular (or interparticle) forces of attraction, the lower the vapor pressure, and the higher the boiling point. The weaker the intermolecular forces of attraction, the higher the vapor pressure, and the lower the boiling point.

Below is a table of the vapor pressures of water at various temperatures. We can see from the table that vapor pressure increases with increasing temperature.

In the figure below, the number of molecules vs kinetic energy is plotted for the same substance at two different temperatures. Again, as the temperature increases, the vapor pressure also increases. Notice how the curve at the higher temperature is broadened and the maximum of the curve is lower than the curve at the lower temperature.

The energy required for the molecules to go into the vapor phase is indicated on the plot. It can be seen that there are more molecules at the higher temperature that have enough energy to go into the vapor phase than there are at the lower temperature.

The boiling point of a liquid is the temperature at which the vapor pressure of the liquid is equal to the external pressure. The normal boiling point of a liquid is observed at an atmospheric pressure of 1 atm (760 mmHg). At the boiling point, a substance is in equilibrium with the liquid and the gas. For example, water has a boiling point of 100^oC. At this temperature there is both liquid water and water vapor.

We all know if we put a paper cup into a flame, the cup will burn. But, if the cup is filled with water, the cup will not burn. This is because all of the thermal energy is being used to heat the water. Once the water has boiled off, then the cup will burn.

Worksheet: Evaporation, Vapor Pressure, and Boiling Point

Exercises

Exercise 1. How does the vapor pressure of a substance change for the following changes?

Check Solutions/Answer to Exercise 1

Exercise 2. Which of the following would have the highest vapor pressure? Explain.

Check Solutions/Answer to Exercise 2

Exercise 3. Which of the following would have the lowest vapor pressure? Explain.

Check Solutions/Answer to Exercise 3

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