The Kinetic Theory of Gases Over time, a single theory of how gases work developed based on the various gas laws. It is known today as the kinetic theory of gases.

All particles are in random, constant, straight-line motion. The particles have kinetic energy. The motion is interrupted by collisions with other molecules, or the container walls – pressure results from these collisions. More collisions = higher pressure

Gas particles (molecules) are separated by great distances relative to their size, so the volume of the gas molecules themselves is insignificant. This explains why gases are easily compressed, and why they mix easily. In other places in the universe, gases are not separated by great distances. Can you name a few places?

The molecules have no attractive forces between them. Again, under high pressure, attraction and repulsion will result in deviations from ideal gas behavior.

Collisions between gas particles may result in transfer of energy between gas particles, but the total energy of the system remains constant. A sample of gas has molecules of different velocities (different KE). The average kinetic energy is proportional to the absolute temperature (Kelvin) of the gas.

Ideal Gases A gas is considered “ideal” if - there are great distances between molecules (low pressure) - there are negligible attractive forces between molecules (high temperature) Under conditions of low pressure and high temperature, most gases behave “ideally”

Examples of Ideal Gases All of the major components of our atmosphere behave ideally at STP except one: Water is not an ideal gas on here on Earth! Water has strong attractive forces compared to N 2, O 2, and Ar Under different conditions, a gas that is ideal here (methane) will be a real gas somewhere else (Titan)

Real Gases There are conditions of temperature and pressure where gases do not behave ideally. Under extremely high pressures, there is little empty space; gases become largely incompressible, like liquids and solids Under low temperatures, molecules are moving slowly and intermolecular forces become a major factor – liquifaction occurs.

PV ≠ k for real gases!

Remember Ideal Gases – high temperatures, low pressures Real Gases – low temperatures, high pressures