General Chemistry/Behaviour of Gases

Evaporation of Liquids

An open bottle of vinegar standing at room temperature will spread odor across the room despite the fact that it is not boiling. This is because kinetic energy is unevenly distributed throughout a liquid. Some molecules are moving faster than others. This allows the liquid to evaporate into a gas. In order for evaporation to occur, the entire liquid does not need to be at the boiling temperature; only some individual molecules. At a given temperature a liquid evaporates at a specific rate, causing gas pressure above the liquid in the system that contains it. Can you guess the vapor pressure of water at 100 degrees Celsius? If you guessed 760 torr or 1 atm, you're right. Liquids boil when their vapor pressures equal the atmospheric pressure around them. Some liquids are more volatile than others, meaning they evaporate more readily.

A Boltzmann distribution at various temperatures.

Evaporation occurs because of the uneven distribution of kinetic energy. This is called a Boltzmann distribution. Although the average kinetic energy of the liquid is below the boiling point, some molecules have above average energy. That gives them enough energy to overcome their bonds and break free of the liquid. At the same time, a gas molecule could strike the liquid and slow down enough to become part of it. There is an equilibrium between the amount of evaporating liquid and the amount of liquifying gas. Higher temperatures will have an equilibrium that favors more gas.

Motion of Gas Particles

This animation demonstrates the motion of gas particles.

Gas particles move in random, straight-line motion. Gas particles possess a greater kinetic energy than the particles of a liquid or solid. The greater kinetic energy is due to gases existing at higher temperatures than liquids or solids. As temperature increases, particles move faster and, thus, have greater kinetic energy. The particles of a gas have minimal interactions, except collisions with each other.

Diffusion and Effusion

Due to their random motion, gases will eventually escape from a container if there is a hole in it. This is called effusion. Gases also spread out across a room, or mix together in a container. This is called diffusion.

More information about diffusion and effusion is presented in the next chapter.

Dalton's Law of Partial Pressures

The total amount of pressure exerted by a gas is equal to the sum of its components:

$P_{total}=P_{1}+P_{2}+P_{3}+...$

For instance, if you add 300 torr of nitrogen gas to a container, then you add 250 torr of oxygen to that container, the total pressure will be 550 torr (assuming volume and temperature stayed constant).

Another way of stating this law is that the pressure exerted by a particular gas in a mixture (gases always form homogeneous mixtures with each other) is equal to its mole fraction multiplied by the total pressure.