1. 1 Torricelli’s Barometer

2. 2 A simple manometer for measuring gas pressure in a container

3. 3 Robert Boyle

4. 4 A J-tube similar to the one used by Boyle

5. 5 Increased pressure leads to decreased volume

6. 6 Table 5.1 Actual Data from Boyle's Experiment

7. 7 Plotting Boyle's Data

8. 8 As pressure increases, the volume of SO2 decreases

9. 9 As pressure increases, the volume decreases

10. 10 Antoine and Marie Lavoisier (Painting by Jacques-Louis David)

11. 11 Empress Eugenie of France (Painting by Franz Winterhalter)

12. 12 Prof. Jacques Charles

13. 13

14. 14

15. 15

16. 16 Increasing the temperature of a gas (at constant pressure) increases its volume.

17. 17

18. 18

19. 19 Plots of V versus T(Celsius) for several gases

20. 20 Plots of V versus T using the Kelvin scale for temperature

21. 21 At constant volume, pressure increases in proportion to Kelvin temperature.

22. 22 Boyle’s law: PV = k (for constant T) Charles’s law: V = kT (for constant P) Gay-Lussac’s law: P = kT (for constant V) COMBINE ALL THREE: PV = k T or PV/T = k for any sample

23. 23 PV = PV T T (for any sample of gas under two sets of conditions)

24. 24

25. 25 One mole of any gas at S.T.P. (273 K, 1.0 atm.) occupies 22.4 L and just fits into this box

26. 26 At a given temperature and pressure, each of these balloons holds the same number of moles.

27. 27 The partial pressure of each gas in a mixture depends on the number of moles of that gas.

28. 28 PV = n RT R = 0.0821 L atm / mol K

29. 29 Kinetic molecular theory models gases as large numbers of randomly moving particles of negligible volume that interact with other particles (and container walls) only by collision.

30. 30 The End

31. 31

32. 32 The production of oxygen by thermal decomposition of KClO3

33. 33 Reaction of zinc with HCl

34. 34 Effusion of a gas into an evacuated chamber

35. 35 Relative molecular speed distribution of H2 and UF6

36. 36 NH3 gas and HCl gas diffuse toward each other and react to form solid NH4Cl

37. 37 Velocity distribution of N2 molecules at 3 different temperatures

38. 38 Slower Molecules Produce a Lower Pressure

39. 39 Gas at low concentration has relatively fewer interactions between particles

40. 40 Pairwise interactions among gas particles

41. 41 Velocity distribution of O2 Molecules at STP

42. 42 The volume taken up by the gas particles themselves is less important (a) at low pressure than (b) at high pressure.

43. 43 Molecular Sieve Model

44. 44 Inflated Air Bags

45. 45 The pressure exerted by the atomsphere can be demonstrated by boiling water in a large metal can

46. 46 Acid Rain: Statue in 1990

47. 47 Schematic diagram of the process for “scrubbing” sulfur dioxide emissions from stack gases in power plants

48. 48 An environmental officer testing the pH of water.

49. 49 Atmospheric composition of dry air near sea level

50. 50 Variation of temperature (blue) and pressure (dashed lines) with altitude

51. 51 Molar volumes for various gases at 0°C and 1 atm

52. 52 Plot of PV versus P for several non-ideal gases at low pressure

53. 53 Plot of PV/nRT versus P for Nitrogen gas at 3 temperatures

54. 54 Plot of PV versus P for 1 mol of ammonia.

55. 55 Plots of PV/nRT versus P for Several Gases (at 200K)

56. 56 Values of the van der Waals constants for selected gases

57. 57 Increased volume due to increased moles of gas at constant temperature and pressure

58. 58 The ratio of the volumes of gaseous N 2 and liquid N 2 is 22.4/0.035=640 and the spacing of the molecules is 9 times farther apart in N 2 (g).

59. 59