AP Chemistry Unit 2

1. AP ChemistryUnit 2 Chapter 5: Pressure; Boyle’s, Charles’, Avogadro’s laws; Ideal gas law; gas stoichiometry; Dalton’s law of partial pressures; KMT of gases; Effusion and diffusion; real gases.

2. CB p. 2 A gas • Uniformly fills any container. • Mixes completely with any other gas • Exerts pressure on its surroundings.

3. CB p. 2 Pressure • is equal to force/unit area • SI units = Newton/meter2 = 1 Pascal (Pa) • 1 standard atmosphere = 1 atm • 1 atm = 760 mm Hg • 1 atm = 760 torr • 1 atm = 101.325 kPa

4. CB p. 2 Figure 5.1a: The pressure exerted by the gases in the atmosphere can be demonstrated by boiling water in a large metal can (a) and then turning off the heat and sealing the can.

5. CB p. 2 Figure 5.01b: As the can cools, the water vapor condenses, lowering the gas pressure inside the can. This causes the can to crumple (b).

6. CB p. 2 Figure 5.2: A torricellian barometer. The tube, completely filled with mercury, is inverted in a dish of mercury.

7. CB p. 2 Figure 5.3: A simple manometer.

8. CB p. 2

9. CB p. 4 Boyle’s law* • Pressure  Volume = Constant (T = constant) P1V1 = P2V2 (T = constant) V 1/P (T = constant) *Holds precisely only at very low pressures.

10. CB p. 4 A gas that strictly obeysBoyle’s Law is called an ideal gas.

11. CB p. 4 Boyle’s law

12. CB p. 4 Figure 5.4: A J-tube similar to the one used by Boyle.

13. CB p. 4 Figure 5.5: Plotting Boyle's data from Table 5.1. (a) A plot of P versus V shows that the volume doubles as the pressure is halved. (b) A plot of V versus 1/P gives a straight line. The slope of this line equals the value of the constant k.

14. CB p. 4 Figure 5.6: A plot of PV versus P for several gases at pressures below 1 atm.

15. CB p. 4 As pressure increases, the volume of SO2 decreases.

16. CB p. 4 Figure 5.7: A plot of PV versus P for 1 mol of ammonia. The dashed line shows the extrapolation of the data to zero pressure to give the "ideal" value of PV of 22.41 L • atm.

17. CB p. 6 Charles’ law The volume of a gas is directly proportional to temperature, and extrapolates to zero at zero Kelvin. V = bT (P = constant) b = a proportionality constant

18. CB p. 6 Figure 5.8: Plots of V versus T (ºC) for several gases.

19. CB p. 6 Figure 5.9a. 5.8 except here the Kelvin scale is used for temperature.

20. CB p. 8

21. CB p. 8 Avogadro’s law For a gas at constant temperature and pressure, the volume is directly proportional to the number of moles of gas (at low pressures). V = an a = proportionality constant V = volume of the gas n = number of moles of gas

22. CB p. 8 Figure 5.10: These balloons each hold 1.0L of gas at 25ºC and 1 atm. Each balloon contains 0.041 mol of gas, or 2.5 x 1022 molecules.

23. CB p. 10 Ideal gas law • An equation of statefor a gas. • “state” is the condition of the gas at a given time. PV = nRT

24. CB p. 10 Ideal gas law PV = nRT R = proportionality constant = 0.08206 L atm  mol P = pressure in atm V = volume in liters n = moles T = temperature in Kelvins Holds closely at P < 1 atm

25. CB p. 10 Standard Temperature and Pressure “STP” P = 1 atmosphere T = C The molar volume of an ideal gas is 22.42 liters at STP

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28. CB p. 10 Figure 5.11: 22.4 L of a gas would just fit into this box.

29. CB p. 12 Dalton’s law of partial pressures For a mixture of gases in a container, PTotal = P1 + P2 + P3 + . . .

30. CB p. 12 Figure 5.12: The partial pressure of each gas in a mixture of gases in a container depends on the number of moles of that gas.

31. CB p. 12

32. CB p. 12 Figure 5.13: The production of oxygen by thermal decomposition of KCIO3. The MnO2 is mixed with the KClO3 to make the reaction faster.

33. CB p. 14

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37. CB p. 16 Kinetic Molecular Theory (KMT) of gases 1. Volume of individual particles is  zero. 2. Collisions of particles with container walls cause pressure exerted by gas. 3. Particles exert no forces on each other. 4. Average kinetic energy  Kelvin temperature of a gas.

38. CB p. 16 The meaning of temperature Kelvin temperature is an index of the random motions of gas particles (higher T means greater motion.)

39. CB p. 16 Figure 5.21: A plot of the relative number of N2 molecules that have a given velocity at three temperatures.

40. CB p. 16

41. CB p. 18 Root-mean-square-velocity • The symbol means the average of the squares of the particle velocities. • The square root of is called the root mean square velocity and is symbolized by u rms.

42. CB p. 18

43. Diffusion: describes the mixing of gases. The rate of diffusion is the rate of gas mixing. Effusion: describes the passage of gas into an evacuated chamber. CB p. 20 Diffusion vs. Effusion

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45. CB p. 20 Figure 5.22: The effusion of a gas into an evacuated chamber.

46. CB p. 20 Figure 5.23: Relative molecular speed distribution of H2 and UF6.

47. CB p. 20 Figure 5.24: (top) When HCl(g) and NH3(g) meet in the tube, a white ring of NH4Cl(s) forms. (bottom) A demonstration of the relative diffusion rates of NH3 and HCl molecules through air.

48. CB p. 22 Real gases • Must correct ideal gas behavior when at high pressure(smaller volume) and low temperature(attractive forces become important).

49. CB p. 22 Real gases   corrected pressure corrected volume Pideal Videal

50. CB p. 22