Chapter 12

1. Chapter 12 Physical Characteristics of Gases

2. Kinetic Molecular Theory • Particles of matter are ALWAYS in motion • Volume of individual particles is  zero. • Collisions of particles with container walls cause pressure exerted by gas. • Particles exert no forces on each other. • Average kinetic energy µ Kelvin temperature of a gas.

3. The Meaning of Temperature • Kelvin temperature is an index of the random motions of gas particles (higher Temp means greater motion & greater ENERGY.)

4. Kinetic Energy of Gas Particles At the same conditions of temperature, all gases have the same average kinetic energy.

5. The Nature of Gases • Gases expand to fill their containers • Gases are fluid – they flow • Gases have low density • 1/1000 the density of the equivalent liquid or solid • Gases are compressible • Gases effuse and diffuse

6. Diffusion Diffusion: describes the mixing of gases. Therateof diffusion is the rate of gas mixing.

7. Effusion Effusion: describes the movement of gas from an area of HIGHER gas concentration to an area of LOWER gas concentration.

8. Pressure • Is caused by the collisions of molecules with the walls of a container • is equal to force/unit area • SI units = Newton/meter2 = 1 Pascal (Pa) • Atmospheres (atm) is the unit most commonly used in Chemistry. • 1 atm = 760 mm Hg = 760 torr = 101,325 Pa = 101.325 kPa

9. Ideal Gases vs Real Gases Ideal gases are imaginary gases that perfectly fit all of the assumptions of the kinetic molecular theory. • Ideal: Gases consist of tiny particles that are far apart relative to their size. • Real: Same! • Collisions between gas particles and between • particles and the walls of the container are • elastic collisions • No kinetic energy is lost in elastic • collisions • Real: No perfectly elastic collisions. Some • energy is lost or gained in every collision.

10. Ideal vs. Real Gases (continued) • Gas particles are in constant, rapid, random motion. • They therefore possess kinetic energy, the energy • of motion • Real: Yes EXCEPT when gas gets VERY cold. • At Absolute zero there is no molecular motion. • Absolute zero = O Kelvein, - 273 oC, or -459 oF • There are no forces of attraction between gas • particles • Real: NO! • All gases are made of protons and electrons • which attract each other! • Polar Gases (H2O & NH3 can have strong • attractions between molecules.

11. Ideal vs. Real Gases (continued) • The kinetic energy of gas particles • depends on temperature, not on the identity • of the particle. • Real: NO. AVERAGEkinetic energy af ALL of • the gas particles depends on the temp. Can not • determine the K. E. of each particle based on • temperature.

12. Measuring Pressure The first device for measuring atmospheric pressure was developed by Evangelista Torricelli during the 17th century. The device was called a “barometer” • Baro = weight • Meter = measure

13. An Early Barometer The normal pressure due to the atmosphere at sea level can support a column of mercury that is 760 mm high.

14. The Aneroid Barometer

15. The Digital Barometer

16. Standard Temperature and Pressure“STP” • P = 1 atmosphere, 760 torr • T = 0°C, 273 Kelvins • The molar volume of an ideal gas is 22.42 liters at STP

17. Robert Boyle(1627-1691) • Boyle was born into an aristocratic Irish family • Became interested in medicine and the new science of Galileo and studied chemistry.  • A founder and an influential fellow of the Royal Society of London • Wrote prolifically on science, philosophy, and theology.

18. Boyle’s Law* Pressure is inversely proportional to volume when temperature is held constant. • Pressure ´ Volume = Constant (k) • P1V1 = P2V2 (T = constant)

19. A Graph of Boyle’s Law

20. Why Don’t I Get a Constant Value for PV = k? • Air is not made • of ideal gases 2. Real gases deviate from ideal behavior at high pressure

21. Jaques Charles (1746-1823) • French Physicist • Conducted the first scientific balloon flight in 1783

22. Charles’s Law • The volume of a gas is directly proportional to temperature, and extrapolates to zero at zero Kelvin. • (P = constant)

23. Converting Celsius to Kelvin Gas law problems involving temperature require that the temperature be in KELVINS! Kelvins = C + 273 °C = Kelvins - 273

24. Joseph Louis Gay-Lussac1778 - 1850 • French chemist and physicist • Known for his studies on the physical properties of gases. • In 1804 he made balloon ascensions to study magnetic forces and to observe the composition and temperature of the air at different altitudes.

25. Gay Lussac’s Law The pressure and temperature of a gas are directly related, provided that the volume remains constant.

26. The Combined Gas Law The combined gas law expresses the relationship between pressure, volume and temperature of a fixed amount of gas.

27. Dalton’s Law of Partial Pressures • For a mixture of gases in a container, • PTotal = P1 + P2 + P3 + . . . This is particularly useful in calculating the pressure of gases collected over water.

28. Standard Molar Volume Equal volumes of all gases at the same temperature and pressure contain the same number of molecules. - Amedeo Avogadro

29. Standard Molar Volume Remember the mole triangle??

30. Ideal Gas Law • PV = nRT • P = pressure in atm • V = volume in liters • T = temperature in Kelvins • n = moles • R = a constant • = 0.0821 L atm/ mol·K Holds closely at P < 1 atm

31. What is R? If PV = nRT and P = atm V = liters n = moles T = Kelvins R = PV SO R = atm x L nT mol x K

32. End of Slide Show