into the gaseous phase. This is called evaporation. If they’re in a closed system
(such as a pot with a lid on it), they quickly lose some kinetic energy and fall
back into the liquid phase, only to be replaced continuously by a couple of other
molecules that manage to escape. They, too, fall back into the liquid phase to be
replaced by other molecules that manage to escape for a few seconds. So a
sample of liquid below its boiling point is always evaporating—a little bit—and
then condensing again (if the liquid is contained). When liquids below their
boiling points are evaporating, a vapor pressure is created. All liquids in a
closed system, at all temperatures, exert some vapor pressure.
But what if the liquid is not in a closed system but is out in the open
environment? Here’s what happens. A little bit evaporates and is blown away.
Then a little more evaporates and is blown or drifts away. Ultimately the whole
sample evaporates. If you put a pot of water outside, even at a temperature of
10°C, it will eventually evaporate, although it will take some time.
Factors Affecting Vapor Pressure
Different liquids differ in their volatility; for instance, if you leave a bucket of
gasoline and a bucket of water outside on a cold day—at a temperature well
below the boiling point of either substance—both will eventually evaporate. But
the gasoline will evaporate much more quickly than the water. This is because
the intermolecular forces that attract gasoline molecules to each other are
weaker than the hydrogen bonds that attract water molecules. Gasoline
molecules need less kinetic energy than water molecules do to overcome the
intermolecular forces that hold them in the liquid state. Because gasoline
evaporates more readily than water, we can say that its vapor pressure is higher,
and it is more volatile than water.
What other factors besides intermolecular force will affect a substance’s vapor
pressure? Well for one, temperature: The higher the temperature, the higher the
average kinetic energy of the molecules. This means more molecules have
enough energy to escape into the gas phase, so there will be more vapor particles
and more vapor pressure above the surface of the liquid. If the container is open
to the environment, the total pressure above the surface of the liquid must equal
atmospheric pressure. The total pressure above the surface is just the vapor
pressure plus the partial pressure of the atmospheric molecules.
atmospheric pressure = vapor pressure +