Chapter 5 Gases

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.

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). (cont’d)

Figure 5. 2: A torricellian barometer Figure 5.2: A torricellian barometer. The tube, completely filled with mercury, is inverted in a dish of mercury.

Figure 5.3: A simple manometer. Copyright © Houghton Mifflin Company. All rights reserved.

Figure 5.4: A J-tube similar to the one used by Boyle.

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Figure 5. 5: Plotting Boyle's data from Table 5. 1 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. Copyright © Houghton Mifflin Company. All rights reserved.

As pressure increases, the volume of SO2 decreases.

Figure 5. 10: These balloons each hold 1. 0L of gas at 25ºC and 1 atm 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.

As pressure increases, the volume decreases.

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Figure 5.11: 22.4 L of a gas would just fit into this box. Copyright © Houghton Mifflin Company. All rights reserved.

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.

Figure 5.14: (a) One mole of N2(l) has a volume of approximately 35 mL and density of 0.81 g/mL. B) One mole of N2(g) has a volume of 22.4 L (STP) and a density of 1.2 x 10-3 g/mL. Thus the ratio of the volumes of gaseous N2 and liquid N2 is 22.4/0.035 = 640 and the spacing of the molecules is 9 times farther apart in N2(g). What about water? Copyright © Houghton Mifflin Company. All rights reserved.

Figure 5.15: The effects of decreasing the volume of a sample of gas at constant temperature. P x V is constant Copyright © Houghton Mifflin Company. All rights reserved.

Figure 5.16: The effects of increasing the temperature of a sample of gas at constant volume. P is proportional to T Copyright © Houghton Mifflin Company. All rights reserved.

Figure 5.17: The effects of increasing the temperature of a sample of gas at constant pressure. V is proportional to T Copyright © Houghton Mifflin Company. All rights reserved.

Figure 5.18: The effects of increasing the number of moles of gas particles at constant temperature and pressure. PV is proportional to n Copyright © Houghton Mifflin Company. All rights reserved.

IDEAL GAS LAW: PV = nRT Copyright © Houghton Mifflin Company. All rights reserved.

Figure 5. 19: Path of one particle in a gas Figure 5.19: Path of one particle in a gas. Any given particle will continuously change its course as a result of collisions with other particles, as well as with the walls of the container. Pressure is molecules colliding with the walls What happens when…. ….you decrease V? ….you increase P? ….you increase n? ….you increase T? Copyright © Houghton Mifflin Company. All rights reserved.

Figure 5.20: A plot of the relative number of O2 molecules that have a given velocity at STP.

Figure 5.21: A plot of the relative number of N2 molecules that have a given velocity at three temperatures

Figure 5.21: A plot of the relative number of N2 molecules that have a given velocity at three temperatures

Figure 5.23: Relative molecular speed distribution of H2 and UF6. Copyright © Houghton Mifflin Company. All rights reserved.

Figure 5.22: The effusion of a gas into an evacuated chamber. Copyright © Houghton Mifflin Company. All rights reserved.

Rate of diffusion. Copyright © Houghton Mifflin Company. All rights reserved.

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. Copyright © Houghton Mifflin Company. All rights reserved.

We can derive PV=nRT from gas kinetics NB: When a gas is not ideal…… Copyright © Houghton Mifflin Company. All rights reserved.

Figure 5.25: Plots of PV/nRT versus P for several gases (200 K). Copyright © Houghton Mifflin Company. All rights reserved.

Figure 5.26: Plots of PV/nRT versus P for nitrogen gas at three temperatures. Copyright © Houghton Mifflin Company. All rights reserved.

Figure 5.27: (a) Gas at low concentration— relatively few interactions between particles. (b) Gas at high concentration—many more interactions between particles. Copyright © Houghton Mifflin Company. All rights reserved.

Figure 5.28: Illustration of pairwise interactions among gas particles.

Van der Waals equation For non-ideal gas Copyright © Houghton Mifflin Company. All rights reserved.

Van der Waals equation For non-ideal gas Corrects for: interaction Excluded volume Copyright © Houghton Mifflin Company. All rights reserved.

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Figure 5.29: The volume taken up by the gas particles themselves is less important at (a) large container volume (low pressure) than at (b) small container volume (high pressure). Copyright © Houghton Mifflin Company. All rights reserved.

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Figure 5.30: The variation of temperature (blue) and pressure (dashed lines) with attitude. Note that the pressure steadily decreases with altitude, but the temperature increases and decreases.

Figure 5.31: Concentration (in molecules per million molecules of "air") for some smog components versus time of day.

Figure 5.33: A schematic diagram of the process for scrubbing sulfur dioxide from stack gases in power plants. Copyright © Houghton Mifflin Company. All rights reserved.