1. Kinetic Theory of Gases Why do gases exert pressure?

2. http://upload.wikimedia.org/wikipedia/comm ons/6/6d/Translational_motion.gif 1. The particles in a gas are in constant, random motion. They are constantly colliding with one another and with the walls of whatever container they’re in. It’s the collisions that cause a gas to exert pressure. 2. The collisions with the container’s walls are perfectly elastic (the particles don’t lose energy when they collide.) If they did, the container would heat up while the gas cooled!

3. 3. The size of the particles is tiny compared to the distances between them. Therefore, the particles don’t interact with one another much (except during collisions). Most of the volume of a gas is really empty space between particles. 1 mole of any gas has the same volume at the same conditions (for example, 22.4 L @ STP) Mixtures of gases in the same container exert pressure on the container’s walls as if each gas were the only gas present. The total pressure is the sum of each gas’s individual pressure.

4. 4. The average kinetic energy of the particles (how hard they impact the walls of the container) depends only on the gas’s temperature. All the different gases together in a single container must be at the same temperature (same kinetic energy) since they’re in contact with one another. KE = ½(mass)(velocity 2 ) = 3 / 2 k T (where k is the Boltzmann constant 1.38 x 10 -23 J/K) Small particles move faster (He atoms move 1370 m/s @ room temp. Xe atoms only 270 m/s.)