Nature of Gases

1. Nature of Gases

2. Properties of Gases • Gases have mass (seems obvious) • Easy to compress gases • Completely fill container they are in

3. Properties of Gases • Different gases can move through each other rapidly (diffusion) • Gases exert pressure • Pressure of a gas depends on its temperature, amount and volume

4. Kinetic Molecular Theory • Sometimes abbreviated KMT • Assumes a gas is made of points without volume that have mass • Most gases are molecules • Noble gases = single atom gases

5. KMT • Particles in a gas must be separated by relatively large distances (can be compressed) • Particles must be in rapid, constant motion (explains diffusion)

6. KMT • Gases exert pressure on containers because their particles collide frequently • Collisions must be perfectly elastic (no energy lost because collisions keep occurring)

7. KMT • Gas pressure increases when temperature increases (and volume is held constant) because temperature is a measure of average kinetic energy of particles • KE = mv2/2

8. KMT • Gas particles exert no force on one another (attractive forces between molecules or atoms are so weak the KMT model assumes the force to be zero)

9. KMT • What are the assumptions of the kinetic molecular theory? • (Hint: the points highlighted, underlined, italicized, and in PURPLE in the previous slides)

10. Measuring Gases • We will deal with 4 variables when it comes to measuring gases • Amount • Volume • Temperature • Pressure

11. Amount of a gas (n) • Measured in moles • 1 mole of ANY gas at STP occupies 22.4 L (molar volume) • Therefore, 1 mole of ANY gas at STP has _________ particles (1 mole of ANYTHING has how many of those things?)

12. Volume • Gas completely fills any container it is in • The volume of the container IS the volume of the gas • Measured in liters, cm3, or mL

13. Temperature (T) • Measured in degrees Celsius, on a practical basis • Celsius is always changed to degrees Kelvin for calculation (WHY?) • 0 °C = ? K • T (K) = T (°C) + 273

14. Pressure (P) • Collisions of gases cause pressure • Gas particles in atmosphere exert pressure • Atmospheric pressure is due to the fact that gas has mass which is attracted by Earth’s gravity

15. Pressure (P) • Force per unit area = pressure • Force in SI = Newton, area = m2 • Pressure in SI = Pascal (Pa) • 1,000 Pascal = 1 kPa

16. Pressure • Another way to measure atmospheric pressure = 1 atm • Pressure exerted by atmosphere at sea level • Torricelli invented way of measuring weight of atmosphere

17. Pressure • Torricelli inverted a glass tube of mercury in a bowl of mercury • Atmospheric pressure pushed the column of Hg inside the tube to a certain height • Unit of pressure = torr

18. Torricelli’s Barometer

19. Pressure • 1 atm = 760 torr (mm of Hg) = 101, 325 Pa = 1013.25 hPa = 1013.25 mbar • Aloha!

20. Now for the laws… • Boyle’s law • Boyle trapped air in a U-shaped tube and measured what happened to the volume when pressure was applied

22. Boyle’s Results • Trial Pressure Volume • 1 1.30 atm 3.10 cm3 • 2 1.60 2.51 • 3 1.90 2.11 • 4 2.20 1.84

23. Boyle’s Results • 1 1.30 atm x 3.10 cm3 = 4.03 • 2 1.60 x 2.51 = 4.016 • 3 1.90 x 2.11 = 4.009 • 4 2.20 x 1.84 = 4.048 • Boyle's law demo

24. Results? • Pressure is proportional to the inverse of the volume • Greater the pressure, _____________ the volume • Practical applications? • Diving • Scuba Bill

25. Mathematical statement • PV = k (at constant temp) • P1V1 = P2V2 • If you know three of the variables, you can calculate the remaining one

26. Charles’s Law • Jacques Charles determined the relationship between temperature and volume • At constant pressure, temperature (in K) is directly proportional to volume • By extrapolation, the idea of “absolute zero” is hinted at

27. Charles’s Law

28. Charles’s Law • If temperature increases and pressure remains constant, volume will ____________. • Practical applications? • Checking tire inflation? When should you do this?

29. Charles’s Law • Mathematical statement • V=kT, so V/T = k • V1T2 = V2T1

30. Gay-Lussac’s Law • Direct relationship between temperature and pressure • P = kT • Aero can pressure

31. Gay-Lussac’s Law • At constant volume, P1 = P2 T1 T2

32. Avogadro’s Law • Equal volumes of gases at same temperature and pressure contain an equal number of particles • All gases exhibit same physical behavior • Larger volume must have greater number of particles

33. Avogadro’s Law • V = kn • From the previous laws, you can note volume changes with temperature and pressure • But for 1 mole at STP, all gases will occupy 22.41 L

34. Dalton’s Law of Partial Pressures • In mixture of gases, each gas exerts a pressure as if it were alone in the container • Reason why air (a mixture of gases) in some ways acts like a single gas

35. Dalton continued • Sum of the partial pressures of all gases in a mixture is equal to the total pressure • PT = pa + pb + pc + …

36. Putting It All Together • Ideal Gas Law combines the work of Boyle, Charles, Gay-Lussac and Avogadro • PV = nRT • Animated gas lab

37. Ideal Gas Law • P = pressure in atm • V = volume in liters • n= number of moles • R = gas constant • (0.0821 atm-L/mol-K) • T = temperature in Kelvins

38. Deviant Behavior • Ideal gas law very useful, but doesn’t apply under certain conditions • Very high pressure (experimental values do not agree with predicted values)

39. Deviant Behavior • Why? • High pressure = gas particles close together, volume of gas molecules (or atoms) becomes significant (remember, we assumed gases are single points of mass with no volume)

40. Deviant Behavior • Low temperature failure of KMT • Gas particles slow down and attractive forces between particles become significant

41. Just for grins…

42. Uh-oh….

43. MLIA

44. Deviant Behavior • DONE • DONE • DONE • DONE • Whatever floats your boat!