1 / 42

Physical Characteristics of Gases

Physical Characteristics of Gases. Chapter 10. Kinetic-molecular theory. Particles of matter are always in motion. Ideal gas. An imaginary gas that perfectly fits all the assumptions of the kinetic-molecular theory. We can often treat real gases as ideal gases and still get good results.

chiko
Download Presentation

Physical Characteristics of Gases

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Physical Characteristics of Gases Chapter 10

  2. Kinetic-molecular theory • Particles of matter are always in motion

  3. Ideal gas • An imaginary gas that perfectly fits all the assumptions of the kinetic-molecular theory. • We can often treat real gases as ideal gases and still get good results.

  4. Assumptions of KMT of gases • Large numbers of tiny particles that are far apart compared to their size • Low density • Easily compressed • Elastic collisions • No kinetic energy is lost when gas particles collide with each other or their container • It can be transferred between particles, but the total kinetic energy remains the same

  5. Gas particles are in continuous, rapid, random motion • There are no attractive or repulsive forces between gas particles • When they hit, they don’t stick together • The average kinetic energy of gas particles depends on the temperature of the gas • Direct relationship

  6. Expansion • Gases have indefinite shape and volume. • They completely fill any container they are in. • They also take the shape of that container. • Because they move rapidly in all directions and don’t stick together.

  7. Fluidity • Gas particles slide past each other. • They can flow • Fluid: something that can flow (can be gas or liquid)

  8. Diffusion • Spontaneous mixing of the particles of two substances caused by their random motion. • Gas particles spread out to fill their new container

  9. Effusion • When gas particles pass through a small opening • Particles leak out of the container

  10. Real gases • Do not completely follow kinetic-molecular theory • Especially deviant at high pressures and low temperatures • Noble gases are closest to ideal • Very polar gases are farthest from ideal

  11. Discuss • Describe the conditions under which a real gas is most likely to behave ideally. • Explain the following properties of gases using the kinetic-molecular theory: expansion, fluidity, low density, compressibility, and diffusion.

  12. Describing gases • Needed: • Volume • Temperature • Number of molecules • Pressure • They are mathematically related.

  13. Pressure • balloon • The force per unit area on a surface.

  14. Force • A push or a pull • Measured in newtons (N). • At the Earth’s surface, 1 kg of mass exerts 9.8 N of force due to gravity.

  15. Pressure of gases • Gases exert pressure on any surface with which they collide. • Depends on volume, temperature, and number of molecules

  16. Atmospheric pressure • Air around Earth exerts a pressure on it’s surface and everything on it. • Like the weight of all the molecules pressing down.

  17. h Barometer • Used to measure atmospheric pressure. • Height of liquid (usually mercury) in tube can be used to express atmospheric pressure. • At sea level, the average is 760 mm Hg.

  18. Manometer • Used to measure the pressure of gases. • The height difference between the two arms is the pressure.

  19. Pressure Units

  20. STP • Standard temperature and pressure. • 0 °C and 1 atm • Used to compare volumes of gases.

  21. Example • A weather report gives a current atmospheric pressure of 745.8 mm Hg. Convert this to • Atmospheres • 0.9813 atm • Torr • 745.8 torr • Kilopascals • 99.43 kPa

  22. Discuss • Define pressure • What is STP? • Convert 151.98 kPa to atmospheres • 1.4999 atm

  23. Boyle’s Law • Fixed: mass and temperature • The volume varies inversely with pressure • Less volume, means the particles hit the walls more often. • This increases the pressure

  24. Boyle’s Law • Mathematically: • Each sample of gas has its own k.

  25. Example • A helium-filled balloon contains 125 mL of gas at a pressure of 0.974 atm. What volume will the gas occupy at standard pressure, assuming constant temperature? • 122 mL

  26. You try • A weather balloon with a volume of 1.375 L is released from Earth’s surface at sea level. What volume will the balloon occupy at an altitude of 20.0 km, where the air pressure is 10.0 kPa, assuming constant temperature? • 13.9 L

  27. Charles’s Law • Fixed: mass and pressure • Volume varies directly with temperature. • As temperature goes up, the particles have more energy, so they hit the walls more often and with more force • This pushes the walls outward.

  28. Charles’s Law • Mathematically

  29. Kelvin Scale • Charles’s law works more elegantly on the Kelvin Scale than the Celsius Scale. • If you double the temperature, the volume doubles. • Not true with Celsius • We must use Kelvin for Charles’s Law.

  30. Kelvin Scale • Absolute zero: lowest possible temperature • All particle motion stops • 0 K, -273.15 °C • Often rounded to 273

  31. Example • A balloon filled with oxygen gas occupies a volume of 5.5 L at 25 °C. What volume will the gas occupy at 100. °C, assuming constant pressure? • 6.9 L

  32. You try • A sample of nitrogen gas is contained in a piston with a freely moving cylinder. At 0.0 °C, the volume of the gas is 375 mL. To what temperature must the gas be heated to occupy a volume of 500. mL, assuming constant pressure? • 91 °C

  33. Gay-Lussac’s Law • Fixed: mass and volume • Pressure varies directly with temperature (in Kelvin) • As temperature goes up, energy of particles goes up. • They go faster and hit the walls harder. • If the walls can’t move, the pressure goes up.

  34. Gay-Lussac’s Law • Mathematically:

  35. You try • The temperature within an automobile tire at the beginning of a long trip is 25 °C. At the conclusion of the trip, the tire has a pressure of 1.80 atm. What is the final Celsius temperature within the tire if its original pressure was 1.75 atm? Assume constant volume. • 34 °C

  36. Combined gas law • Expresses the relationship between pressure, volume, and temperature of a fixed amount of gas.

  37. You try • The volume of a gas at 27.0 °C and 0.200 atm is 80.0 mL. What volume will the same gas sample occupy at standard conditions? • 14.6 mL

  38. Dalton’s Law • The total pressure in a container is the sum of the partial pressures of all the gases in the container.

  39. Application • We can collect gases by displacing water. • When we do this, • Read Patm from the barometer. Look up Pwater in table A-8 in the appendix.

  40. Example • A student has stored 100.0 mL of neon gas over water on a day when the temperature is 27.0 °C. If the barometer in the room reads 743.3 mm Hg, what is the pressure of the neon gas in its container? • 716.6 mm Hg

  41. You try • A sample of nitrogen gas is collected over water at a temperature of 23.0 °C. What is the pressure of the nitrogen gas if atmospheric pressure is 785 mm Hg? • 764 mm Hg

More Related