1. Gases The Kinetic-Molecular Theory

2. Ludwig Boltzman and James Maxwell
They each proposed a model to explain the properties of gases.
This model is called, kinetic-molecular theory because all of the gases known to them contained molecules.
The word kinetic means “to move” and objects in motion have kinetic energy.

3. Ludwig Boltzman James Maxwell

4. The Kinetic-Molecular Theory
Describes the behavior of gases in terms of particles in motion.
This model makes several assumptions about the size, motion, and energy of gas particles.

5. Particle Size
Gases consist of small particles that are separated from one another by empty space.
Gas particles are far apart, so there is no significant attractive or repulsive forces among them.

6. Particle Motion Gas particles are in constant random motion.
Particles move in a straight line until they collide with other particles or with the walls of their container.
Collisions between gas particles are elastic, meaning that no kinetic energy is lost.
Kinetic energy can be transferred between colliding particles, but the total energy of the particles does not change.

7. Particle Energy
2 factors determine the kinetic energy of a particle. Mass and Velocity
All particles do not have the same velocity but have the same mass.

8. Explaining the Behavior of Gases

9. Low Density Density is mass per unit volume
The kinetic-molecular theory states a great deal of space exists between gas particles

10. Compression and Expansion
When a gas is compressed the particles get closer
When a gas expands there is more air space

11. Diffusion and Effusion
There are no significant attractions between gas particles
The mixture of gases in the air diffuse until they are evenly distributed
The rate of diffusion depends mainly on the mass of particles involved
Mass of a gas varies from gas to gas

12. Diffusion and Effusion cont.
During effusion a gas escapes through a tiny opening
Thomas Graham did experiments to measure the rates of effusion for different gases at the same temperature

13. Diffusion and Effusion cont.
Grahams law of effusion states that the rate of effusion for a gas is inversely proportional to the square root of its molar mass
Grahams law also applies to rates of diffusion

14. GAS PRESSURE PRESSURE: FORCE PER UNIT AREA
EX: Water Striders, snowshoes

15. MEASURING AIR PRESSURE
TOOLS:
BAROMETER: MEASURES ATMOSPHERIC PRESSURE
MANOMETER: MEASURES THE PRESSURE OF A GAS IN A CLOSED CONTAINER

16. UNITS OF PRESSURE Pascal: N/m^2 (SI unit of pressure)
COMPARED TO ATM
KILOPASCAL (kPa)
1 atm = kPa
MILLIMETERS OF MERCURY (mm Hg)
1 atm = 760 mm Hg
TORR
1 atm = 760 torr
POUNDS PER SQUARE INCH (lb./in^2) or (psi)
1 atm = 14.7 psi
Atmosphere (atm)
Pascal: N/m^2 (SI unit of pressure)
Atmosphere: 760 mm Hg
@ sea 0 degrees C

17. Dalton’s Law of Partial Pressures
The total pressure of a mixture of gases is equal to the sum of the pressures of all the gases in the mixture.
Ex: P (gas 1) + P (gas 2) = P (total)