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Gases The Kinetic-Molecular Theory

Gases The Kinetic-Molecular Theory. 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.

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Gases The Kinetic-Molecular Theory

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  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) • Atmosphere: 760 mm Hg • @ sea level, @ 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)

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