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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,

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Presentation on theme: "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,"— Presentation transcript:

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 BoltzmanJames 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 UnitCOMPARED TO ATM KILOPASCAL (kPa)1 atm = 101.3 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 level, @ 0 degrees C

17 Daltons 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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