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Physical Characteristics of Gases

Physical Characteristics of Gases. Chapter 10. Section 1:. The Kinetic-Molecular Theory of Matter. Section 1: Kinetic-Molecular Theory of Matter. Kinetic-molecular theory is based on… All matter is made of particles that are in constant motion.

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Physical Characteristics of Gases

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  1. Physical Characteristics of Gases Chapter 10

  2. Section 1: The Kinetic-Molecular Theory of Matter

  3. Section 1: Kinetic-Molecular Theory of Matter • Kinetic-molecular theory is based on… All matter is made of particles that are in constant motion. • Is used to explain properties of solids, liquids, & gases.

  4. Section 1: Kinetic-Molecular Theory of Matter • Solids, liquids, and gases vary due to the energy of the particles and the forces that act upon them. • This chapter will study gases.

  5. Section 1: Kinetic-Molecular Theory of Matter • Theory pertains to idealgases - notrealgases. • An ideal gas is an imaginary gas that perfectly fits all the assumptions of the k-m theory.

  6. Section 1: Kinetic-Molecular Theory of Matter 5 assumptions: 1.Gases consist of large numbers of tiny particles that are far apart relative to their size • Much farther apart than liquids & solids so they can be compressed.

  7. Section 1: Kinetic-Molecular Theory of Matter 5 assumptions (continued): 2.Collisions are elastic - there is no net loss of kinetic energy. Kinetic energy is completely transferred during collisions.

  8. Section 1: Kinetic-Molecular Theory of Matter 5 assumptions (continued): Kinetic energy is constant (if at same temperature)

  9. Section 1: Kinetic-Molecular Theory of Matter 5 assumptions (continued): 3.Gas particles are in constant, rapid, random motion. They have kinetic energy.

  10. Section 1: Kinetic-Molecular Theory of Matter 5 assumptions (continued): 4. No forces of attraction act on gas particles.

  11. Section 1: Kinetic-Molecular Theory of Matter 5 assumptions (continued): • The average kinetic energy of gas particles depends on the temperature of the gas. KE= mv2 2

  12. Section 1: Kinetic-Molecular Theory of Matter Consider: KE= mv2 2 What does KE depend on if gases are the same kind? What does KE depend on if gases are at same temperature but are differentkinds of gases?

  13. Physical Properties of Gases • Expansion – Completely fill container & take its shape

  14. Physical Properties of Gases • Fluidity – Gases have the ability to flow. Particles can glide past each other.

  15. Physical Properties of Gases • Low density: Gas densities are about 1/1000 that of the liquid or solid phase. • Compressibility: Steel canisters contain about 100 times the number of particles than at normal pressure.

  16. Physical Properties of Gases • Diffusion: Gases randomly move and mix by random motion of their particles. • The rate of diffusion depends on the mass of the particles. Heavier ones move more slowly.

  17. Animation of diffusion

  18. Physical Properties of Gases • Effusion: Gases under pressure spread out when released from a small opening.

  19. Deviations of Real Gases • Real gases do not behave according to all the assumptions of the K-M theory. • Real gases deviate most when they are under very high pressures and very cold temperatures.

  20. Reason behind this… • Both high pressure and colder temperatures force the atoms or molecules of a gas closer together. • When real gases get closer together they experience intermolecular attractions and then they condense to form liquids.

  21. Deviations of Real Gases • Noble gases behave more like ideal gases than any others. • More polar gases behave less like ideal gases.

  22. Section 2: Pressure

  23. Section 2: Pressure To describe a gas you must state 4 quantities: • Volume • Temperature • Number of molecules • Pressure

  24. Section 2: Pressure Pressure is defined as the amount of force per unit of area. P = force area

  25. Section 2: Pressure P = force area Force unit is Newtons Area unit is square meters Pressure = N/m2

  26. Section 2: Pressure Open can Closed Can Air Pumped Out

  27. Measuring Pressure • What tool do we use to measure atmospheric pressure? • Barometer! • First built by Evangelista Torricelli (1600’s)

  28. Measuring Pressure • Torricelli noticed that pumps could raise water only 34 feet high. • He compared density of mercury to density of water (14x greater) • Predicted height that mercury could be raised (1/14 of 34 ft or about 30 inches).

  29. Units of Pressure • Several different units are used for pressure: • inches of mercury (ex. 30.4 and rising) • mm of Hg • atm (atmospheres) • torrs • Kilopascals (1 Kpa= 1N/m2)

  30. Units of Pressure • Standard pressure taken at sea level: • 29.9 inches of mercury • 760 mm of Hg • 1 atm (atmospheres) • 760 torrs • 101.3 Kilopascals

  31. Section 3 The Gas Laws Mathematical relationships between volume, temperature, pressure and quantity of gases

  32. Section 3: The Gas Laws • Boyle’s Law • Charles’ Law • Dalton’s Law of Partial Pressures • Gay-Lussac’s Law • Combined Gas Law

  33. Section 3: The Gas Laws • Boyle’s Law • Relates gas volume to pressure • Has an inverse relationship P ↑ V ↓

  34. Section 3: The Gas Laws • Formula for Boyle’s Law P1V1 = P2V2 What would be a gas’s volume if the pressure reduced from 98 kPa down to 60 kPa if its original volume was 300 Liters?

  35. Simulation of Boyle's Law

  36. Section 3: The Gas Laws P1V1 = P2V2 98 (300) = 60 (V2) V2 = 490 L

  37. Section 3: The Gas Laws Charles’ Law Relates temperature to gas volume. Directly proportional. Is based on absolute zero.

  38. Section 3: The Gas Laws In 1787, Jacques Charles found that as he decreased the temperature of a gas 1 degree then the volume decreased 1/273.

  39. Section 3: The Gas Laws The volume of a fixed mass of gas at constant pressure varies directly with the Kelvin temperature.

  40. Section 3: The Gas Laws Formula for Charles’ Law (temperature must be in Kelvin) V1 V2 = T1 T2

  41. Charles’ Law Graph V T

  42. Problem: • If the temperature of a gas increases from 25 degrees Celsius up to 80 degree Celsius and the original volume was 10 liters of gas, then what would the final volume be???? 10 X (25 + 273) (80 + 273) X = 11.8 liters

  43. Gay-Lussac’s Law • Relates pressure and temperature (assumes constant volume) P1 P2 = T1 T2

  44. Combined Gas Law • This law is used when 2 variables change – pressure, temperature, or volume. P1V1 = P2V2 T1 T2

  45. Dalton’s Law of Partial Pressure • If there is no chemical reaction occurring then the pressure of a mixture of gases will be equal to all the combined pressures of each gas. • Equation: PT = P1 + P2 + P3 + …

  46. Collecting Gases by Water Displacement • Water is commonly collected by bubbling it through water. • The pressure then in the container is a combination of both the pressure of the gas as well as a small amount of water vapor.

  47. Gas Collection

  48. Collecting Gases by Water Displacement • To determine the pressure of the gas you must subtract the pressure of the water vapor. • The pressure of the water varies according to the temperature. • Use the chart for H2O pressure:

  49. Vapor Pressure of Water

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