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Gases Ch. 6

Gases Ch. 6. Chemistry II Milbank High School. Kinetic Molecular Theory. All matter is composed of tiny, discrete particles called molecules They are in rapid, constant motion Move in straight lines Molecules are small compared to the spaces between them Little attraction between molecules

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Gases Ch. 6

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  1. GasesCh. 6 Chemistry II Milbank High School

  2. Kinetic Molecular Theory • All matter is composed of tiny, discrete particles called molecules • They are in rapid, constant motion • Move in straight lines • Molecules are small compared to the spaces between them • Little attraction between molecules • Molecules collide with one another, and energy is conserved • Temperature is the measure of the average kinetic energy of gas molecules

  3. Kinetic Molecular Theory Con’t • Little attraction between molecules • Molecules collide with one another, and energy is conserved • Temperature is the measure of the average kinetic energy of gas molecules

  4. Properties of Gases • Compressible • Low density • Don’t settle due to gravity (other than the density of air) • Hot particles move faster

  5. Atmospheric Pressure • Pressure is force per unit area P = F/A • SI unit for pressure is the pascal (Pa 1 Pa = 1 N / m2 • Barometer: used to measure pressure

  6. Pressure Units • 1 atm = 760 mmHg = 760 torr = 101.3 kPa = 29.921 in.Hg = 14.696 lb/in2

  7. GAS LAWS

  8. Boyle’s Law • Volume of a gas varies inversely with pressure • PV = PV • Reason gases can be stored in a small volume (high pressure)

  9. Boyle’s Law and Breathing • Inspiration (breathing in)—diaphragm lowered and chest is expanded, increasing the volume of the chest cavity • Increased volume = decreased pressure so air enters from a high pressure area (outside air) to low pressure (lungs)

  10. Charles’s Law • Volume is directly proportional to Kelvin temperature • V/T = V/T

  11. Avogadro’s Law • The volume of a gas is directly proportional to the number of gas molecules present • STP (273 K and 1 atm) • 1 mol of gas = 22.4 L

  12. Combined Gas Law • PV/T = PV/T Just make sure units are the same

  13. Ideal Gas Law • Takes into account different amounts of gases • PV = nRT • n is # of mol • R = . 0821 L·atm/mol·K (universal gas constant)

  14. Henry’s Law • The solubility of a gas in a liquid at a given temp is directly proportional to the pressure of the gas at the surface of the liquid • Pop bottle: bottle is under pressure, when opened the pressure is reduced thus the solubility is reduced (bubbles of CO2 escape)

  15. Henry’s Law and Deep Sea Diving • Divers need compressed air • Compressed air is much more soluble in blood • N2 acts as a narcotic at depths below 30m because of the increased pressure (Nitrogen narcosis) • Decompression sickness (the “bends”) • Bubbles of nitrogen gas in blood and other tissues • Occurs when there is a sudden drop in atmospheric pressure • Diver, miners, passengers in airplane

  16. Dalton’s Law of Partial Pressures • Total pressure of the gases in a mixture is equal to the sum of the partial pressures of the separate gases Ptotal = P1 + P2 +P3 + ….. • Gases collected over water • Thus contain water vapor • Vapor pressure: partial pressure exerted by its molecules in the gas phase above the liquid phase

  17. Dalton’s Law of Partial Pressures Con’t • Humidity: measure of water vapor in air • Relative humidity: actual amount of water vapor compared to amount the air could hold at the temperature • Cool air holds less, thus as temp falls at night, water condenses from the air to form dew

  18. Daltons’ Law of Partial Pressures Con’t • Respiration • Inspired air: high in O2, low in CO2 • Cellular fluid: low in O2, high in CO2 • Gases transferred through diffusion • Gases flow from an area of high concentration to an area of low concentration

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