Vapor Pressure, Evaporation and Boiling

Vapor Pressure, Evaporation and Boiling

Part 1:

Pressure

All gases exert some amount of pressure on their surroundings. This pressure comes from the fact that gas molecules are constantly moving, at random speeds and directions.

The amount of pressure exerted by the gas depends on two things: How fast the molecules are moving, or their temperature, and how many molecules there are in a certain amount of space.

If you take a certain volume of gas and you increase its temperature, the molecules will move faster, and therefore will hit the sides of the container with more force, thereby increasing its pressure. If the container can change its shape (like a balloon) then the volume will increase to relieve the pressure.

If you take a certain volume of a gas and you squeeze it to shrink its volume, its pressure goes up. This is because the molecules will be hitting the side of the container more forcefully as you try to squeeze them together.

Atmospheric Pressure

We live in a “sea” of air that surrounds the Earth – this is called the atmosphere. This air is pulled down to the Earth by the Earth’s gravity. You can think of the atmosphere as being squeezed down to the Earth – just like a balloon that is being squeezed. Therefore, there is a tremendous pressure being exerted on all objects inside the atmosphere. The higher up you go (such as to the top of a mountain) the amount of air pushing down on you decreases. Therefore the pressure goes down the farther you go up away from the surface of the Earth.

Vapor Pressure and Evaporation

Everyone knows that an open container of water (such as a glass of water) will evaporate over time, even at room temperature, but we probably don't give this observation much thought.

The term "vapor" is another word for the gas state of anything. Water, gasoline, rubbing alcohol, and finger nail polish remover (ethyl acetate) are all normally liquids at room temperature, but they all evaporate to give a gas, or vapor.

To begin to see what’s going on, think about a glass of water. In the picture below, we can see that the molecules on the surface are in direct contact with the air. These molecules have only about half as many neighbors as do molecules inside the liquid. Therefore, the forces holding them in the liquid are slightly lower than molecules in the rest of the liquid. (The surface area determines the rate at which a given volume of liquid will evaporate. A flat pan of water will evaporate more quickly than a narrow glass, even for the same volume.)

The molecules escaping from the liquid, because they become a gas, exert a certain pressure. This pressure is called the “vapor pressure”. The vapor pressure is, in a sense, a measure of the number of molecules becoming a gas.

Questions for Part 1:

1. What causes gases to exert a pressure on their surroundings?
2. Draw your “mental model” for what happens to the pressure on a container holding a gas (like a balloon) when the temperature goes up.
3. Draw your “mental model” for what happens to the pressure on a container holding a gas (like a balloon) when it is squeezed.
4. Why does the Earth’s atmosphere exert a pressure on us?
5. What does the term “vapor” mean?
6. What does the term “vapor pressure” mean?
7. Why will molecules at the top of a liquid evaporate more easily than those deeper in the liquid?

Part 2:

Evaporation and Energy

When molecules are evaporating, they must do two things: (a) get fast enough to become molecules in the gas state, and (b) gain enough energy to prevent themselves from falling back into the liquid. Both of these take energy. Where does this extra energy come from? It comes from the surrounding atmosphere or from the surrounding water molecules. They give up some energy when they collide with the molecules that evaporate.

All the molecules in the liquid are all moving randomly at different speeds. Some of them are moving very slowly, and some are moving quickly. Some of them are moving fast enough to have the energy to break free of the forces holding them together. These molecules will break free and move into the vapor, or gas state.

The number of molecules with enough energy to break free is proportional to the temperature. Because temperature is a measure of the average speed of the molecules, as the temperature is raised, more molecules have enough energy to break free and become a gas. Therefore, the vapor pressure increases with increasing temperature.

Effect of Vapor Pressure on Boiling

When the temperature of a liquid increases and gets closer to the boiling point, more and more molecules gain the energy necessary to evaporate. This increases the Vapor Pressure. In general, what keeps most of the molecules from escaping from the liquid is the fact that most of them are being pushed back down into the liquid by the atmospheric pressure (atmospheric pressure is the pressure exerted by the air around us). If you increase the vapor pressure enough, by raising the temperature, the pressure of molecules escaping the liquid exceeds the pressure pushing back down on them. At this point the Vapor Pressure exceeds Atmospheric Pressure and Boiling occurs. In fact the definition of boiling is the point at which the vapor pressure of a liquid is exceeds the atmospheric pressure.

Questions for Part 2:

1. In order for a molecule to escape into the gas state, what 2 things must be true?
2. Your friend tells you that temperature measures the speed of molecules in a substance. He also tells you that at a certain temperature, all the molecules must be moving at the same speed. What’s wrong with your friend’s ideas?
3. What happens to the vapor pressure at increased temperature?
4. In your own words, what is the definition of boiling point?
5. Explain, using the ideas in this article, why water can boil at room temperature if you lower the atmospheric pressure enough.
6. Draw your “mental model” to help explain #5.