# Sample Exercise 10.1 Converting Pressure Units

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1 Sample Exercise 10.1 Converting Pressure Units (a) Convert atm to torr. (b) Convert torr to atmospheres. (c) Convert kpa to torr. Solution Analyze In each case we are given the pressure in one unit and asked to convert it to another unit. Our task, therefore, is to choose the appropriate conversion factors. Plan We can use dimensional analysis to perform the desired conversions. Solve (a) To convert atmospheres to torr, we use the relationship 760 torr = 1 atm: Note that the units cancel in the required manner. (b) We use the same relationship as in part (a). To get the appropriate units to cancel, we must use the conversion factor as follows: (c) The relationship 760 torr = kpa allows us to write an appropriate conversion factor for this problem: 760 torr =

2 Sample Exercise 10.1 Converting Pressure Units Check In each case, compare the magnitude of the answer with the starting value. The torr is a much smaller unit than the atmosphere (since there are 760 torr in 1 atm), so we expect the numerical answer to be larger than the starting quantity in (a) and smaller in (b). In (c) notice that there are nearly 8 torr per kpa, so the numerical answer in torr should be about eight times larger than its value in kpa, consistent with our calculation. Practice Exercise (a) In countries that use the metric system, atmospheric pressure in weather reports is given in kilopascals. Convert a pressure of 745 torr to kilopascals. (b) The pressure at the center of Hurricane Katrina was 902 mbar (millibars). There are 1000 mbar in 1 bar; convert this pressure to atmospheres. Answers: (a) 99.3 kpa, (b) atm

3 Sample Exercise 10.2 Using a Manometer to Measure Gas Pressure On a certain day a laboratory barometer indicates that the atmospheric pressure is torr. A sample of gas is placed in a flask attached to an open-end mercury manometer ( FIGURE 10.3), and a meter stick is used to measure the height of the mercury in the two arms of the U tube. The height of the mercury in the open-end arm is mm, and the height in the arm in contact with the gas in the flask is mm. What is the pressure of the gas in the flask (a) in atmospheres, (b) in kilopascals? Solution Analyze We are given the atmospheric pressure (764.7 torr) and the mercury heights in the two arms of the manometer and asked to determine the gas pressure in the flask. Recall that millimeters of mercury is a pressure unit. We know that the gas pressure from the flask must be greater than atmospheric pressure because the mercury level in the arm on the flask side (103.8 mm) is lower than the level in the arm open to the atmosphere (136.4 mm). Therefore, the gas from the flask is pushing mercury from the arm in contact with the flask into the arm open to the atmosphere. Plan We will use the difference in height between the two arms (h in Figure 10.3) to obtain the amount by which the pressure of the gas exceeds atmospheric pressure. Because an open-end mercury manometer is used, the height difference directly measures the pressure difference in mm Hg or torr between the gas and the atmosphere.

4 Sample Exercise 10.2 Using a Manometer to Measure Gas Pressure Solve (a) The pressure of the gas equals the atmospheric pressure plus h: We convert the pressure of the gas to atmospheres: (b) To calculate the pressure in kpa, we employ the conversion factor Between atmospheres and kpa: Check The calculated pressure is a bit more than 1 atm, which is about 101 kpa. This makes sense because we anticipated that the pressure in the flask would be greater than the atmospheric pressure (764.7 torr = 1.01 atm) acting on the manometer. Practice Exercise Convert a pressure of atm into Pa and kpa. Answers: Pa and 98.8 kpa

5 Sample Exercise 10.3 Evaluating the Effects of Changes in P, V, n, and T on a Gas Suppose we have a gas confined to a cylinder with a movable piston. (Sections 5.2, 5.3) Consider the following changes (assuming no leaks): (a) Heat the gas from 298 K to 360 K at constant pressure. (b) Reduce the volume from 1 L to 0.5 L at constant temperature. (c) Inject additional gas, keeping temperature and volume constant. Indicate how each change affects the average distance between molecules, the pressure of the gas, and the number of moles of gas in the cylinder. Solution Analyze We need to think how each change affects (1) the distance between molecules, (2) the pressure of the gas, and (3) the number of moles of gas in the cylinder. Plan We will use the gas laws and the general properties of gases to analyze each situation. Solve (a) Heating the gas while maintaining constant pressure will cause the piston to move and the volume to increase (Charles s law). Thus, the distance between molecules will increase. At constant pressure, obviously, the pressure will not change. The total number of moles of gas remains the same. (b) Compressing the gas into a smaller volume does not change the total number of gas molecules; thus, the total number of moles remains the same. The average distance between molecules, however, must decrease because of the smaller volume. The reduction in volume causes the pressure to increase (Boyle s law).

6 Sample Exercise 10.3 Evaluating the Effects of Changes in P, V, n, and T on a Gas (c) Injecting more gas into the cylinder while keeping the volume and temperature constant results in more molecules and, thus, an increase in the number of moles of gas in the cylinder. The average distance between molecules must decrease because their number per unit volume increases. Avogadro s law tells us that the volume of the cylinder should have increased when we added more gas, but here the volume is fixed. Boyle s law comes to our aid: If the volume is low, then pressure is high. Therefore, we expect that the pressure will increase in the cylinder if we inject more gas, keeping volume and temperature constant. Practice Exercise Recall that density is mass per volume. (Section 1.4) What happens to the density of a gas as (a) the gas is heated in a constant-volume container; (b) the gas is compressed at constant temperature; (c) additional gas is added to a constant-volume container? Answers: (a) no change, (b) increases, (c) increases

7 Sample Exercise 10.4 Using the Ideal-Gas Equation Calcium carbonate, CaCO 3 (s), the principal compound in limestone, decomposes upon heating to CaO(s) and CO 2 (g). A sample of CaCO 3 is decomposed, and the carbon dioxide is collected in a 250-mL flask. After decomposition is complete, the gas has a pressure of 1.3 atm at a temperature of 31 C. How many moles of CO 2 gas were generated? Solution Analyze We are given the volume (250 ml), pressure (1.3 atm), and temperature (31 C) of a sample of CO 2 gas and asked to calculate the number of moles of CO 2 in the sample. Solve In analyzing and solving gas law problems, it is helpful to tabulate the information given in the problems and then to convert the values to units that are consistent with those for R ( L-atm/mol-K). In this case the given values are Remember: Absolute temperature must always be used when the ideal-gas equation is solved. We now rearrange the ideal-gas equation (Equation 10.5) to solve for n Plan Because we are given V, P, and T, we can solve the ideal-gas equation for the unknown quantity, n.

8 Sample Exercise 10.4 Using the Ideal-Gas Equation Check Appropriate units cancel, thus ensuring that we have properly rearranged the ideal-gas equation and have converted to the correct units. Practice Exercise Tennis balls are usually filled with either air or N 2 gas to a pressure above atmospheric pressure to increase their bounce. If a tennis ball has a volume of 144 cm 3 and contains 0.33 g of N 2 gas, what is the pressure inside the ball at 24 C? Answer: 2.0 atm

9 Sample Exercise 10.5 Calculating the Effect of Temperature Changes on Pressure The gas pressure in an aerosol can is 1.5 atm at 25 C. Assuming that the gas obeys the ideal-gas equation, what is the pressure when the can is heated to 450 C? Solution Analyze We are given the initial pressure (1.5 atm) and temperature (25 C) of the gas and asked for the pressure at a higher temperature (450 C). Plan The volume and number of moles of gas do not change, so we must use a relationship connecting pressure and temperature. Converting temperature to the Kelvin scale and tabulating the given information, we have Solve To determine how P and T are related, we start with the ideal-gas equation and isolate the quantities that do not change (n, V, and R) on one side and the Variables (P and T) on the other side. Because the quotient P/T is a constant, we can write (where the subscripts 1 and 2 represent the initial and final states, respectively). Rearranging to solve for P 2 and substituting the given data give

10 Sample Exercise 10.5 Calculating the Effect of Temperature Changes on Pressure Check This answer is intuitively reasonable increasing the temperature of a gas increases its pressure. Comment It is evident from this example why aerosol cans carry a warning not to incinerate. Practice Exercise The pressure in a natural-gas tank is maintained at 2.20 atm. On a day when the temperature is 15 C, the volume of gas in the tank is m 3. What is the volume of the same quantity of gas on a day when the temperature is 31 C? Answer: m 3

11 Sample Exercise 10.6 Calculating the Effect of Changing P and T on Gas Volume An inflated balloon has a volume of 6.0 L at sea level (1.0 atm) and is allowed to ascend until the pressure is 0.45 atm. During ascent, the temperature of the gas falls from 22 C to 21 C. Calculate the volume of the balloon at its final altitude. Solution Analyze We need to determine a new volume for a gas sample when both pressure and temperature change. Plan Let s again proceed by converting temperatures to kelvins and tabulating our information. Because n is constant, we can use Equation Solve Rearranging Equation 10.8 to solve for V 2 gives

12 Sample Exercise 10.6 Calculating the Effect of Changing P and T on Gas Volume Check The result appears reasonable. Notice that the calculation involves multiplying the initial volume by a ratio of pressures and a ratio of temperatures. Intuitively, we expect decreasing pressure to cause the volume to increase. Similarly, decreasing temperature should cause the volume to decrease. Because the pressure difference is more dramatic than the temperature difference, we expect the effect of the pressure change to predominate in determining the final volume, as it does. Practice Exercise A 0.50-mol sample of oxygen gas is confined at 0 C and 1.0 atm in a cylinder with a movable piston. The piston compresses the gas so that the final volume is half the initial volume and the final pressure is 2.2 atm. What is the final temperature of the gas in degrees Celsius? Answer: 27 C

13 Sample Exercise 10.7 Calculating Gas Density What is the density of carbon tetrachloride vapor at 714 torr and 125 C? Solution Analyze We are asked to calculate the density of a gas given its name, its pressure, and its temperature. From the name we can write the chemical formula of the substance and determine its molar mass. Plan We can use Equation to calculate the density. Before we can do that, however, we must convert the given quantities to the appropriate units, degrees Celsius to kelvins and pressure to atmospheres. We must also calculate the molar mass of CCl 4. Solve The absolute temperature is = 398 K. The pressure is (714 torr) (1 atm/760 torr) = The molar mass of CCl 4 is (4) (35.45) = g/mol. Therefore, Check If we divide molar mass (g/mol) by density (g/l), we end up with L/mol. The numerical value is roughly 154/4.4 = 35. That is in the right ballpark for the molar volume of a gas heated to 125 C at near atmospheric pressure, so our answer is reasonable.

14 Sample Exercise 10.7 Calculating Gas Density Practice Exercise The mean molar mass of the atmosphere at the surface of Titan, Saturn s largest moon, is 28.6 g/mo. The surface temperature is 95 K, and the pressure is 1.6 atm. Assuming ideal behavior, calculate the density of Titan s atmosphere. Answer: 5.9 g/l

15 Sample Exercise 10.8 Calculating the Molar Mass of a Gas A large evacuated flask initially has a mass of g. When the flask is filled with a gas of unknown molar mass to a pressure of 735 torr at 31 C, its mass is g. When the flask is evacuated again Solution Analyze We are given the temperature (31 C) and pressure (735 torr) for a gas, together with information to determine its volume and mass, and we are asked to calculate its molar mass. Solve The gas mass is the difference between the mass of the flask filled with gas and the mass of the evacuated flask: The gas volume equals the volume of water the flask can hold, calculated from the mass and density of the water. The mass of the water is the difference between the masses of the full and evacuated flask: and then filled with water at 31 C, its mass is g. (The density of water at this temperature is g/ml.) Assuming the ideal-gas equation applies, calculate the molar mass of the gas. Plan We need to use the mass information given to calculate the volume of the container and the mass of the gas in it. From this we calculate the gas density and then apply Equation to calculate the molar mass of the gas g g = g g g = g Rearranging the equation for density (d = m/v), we have

16 Sample Exercise 10.8 Calculating the Molar Mass of a Gas Knowing the mass of the gas (2.889 g) and its volume (0.936 L), we can calculate the density of the gas: g/0.936 L = 3.09 g/l After converting pressure to atmospheres and temperature to kelvins, we can use Equation to calculate the molar mass: Check The units work out appropriately, and the value of molar mass obtained is reasonable for a substance that is gaseous near room temperature. Practice Exercise Calculate the average molar mass of dry air if it has a density of 1.17 g/l at 21 C and torr. Answer: 29.0 g/mol

17 Sample Exercise 10.9 Relating a Gas Volume to the Amount of Another Substance in a Reaction Automobile air bags are inflated by nitrogen gas generated by the rapid decomposition of sodium azide, NaN 3 : 2 NaN 3 (s) 2 Na(s) + 3 N 2 (g) If an air bag has a volume of 36 L and is to be filled with nitrogen gas at 1.15 atm and 26.0 C, how many grams of NaN 3 must be decomposed? Solution Analyze This is a multistep problem. We are given the volume, pressure, and temperature of the N 2 gas and the chemical equation for the reaction by which the N 2 is generated. We must use this information to calculate the number of grams of NaN 3 needed to obtain the necessary N 2. Plan We need to use the gas data (P, V, and T) and the ideal-gas equation to calculate the number of moles of N 2 gas that should be formed for the air bag to operate correctly. We can then use the balanced equation to determine the number of moles of NaN 3 needed. Finally, we can convert moles of NaN 3 to grams.

18 Sample Exercise 10.9 Relating a Gas Volume to the Amount of Another Substance in a Reaction Solve The number of moles of N 2 is determined using the ideal-gas equation: We use the coefficients in the balanced equation to calculate the number of moles of NaN 3 : Finally, using the molar mass of NaN 3, we convert moles of NaN 3 to grams: Check The units cancel properly at each step in the calculation, leaving us with the correct units in the answer, g NaN 3. Practice Exercise In the first step in the industrial process for making nitric acid, ammonia reacts with oxygen in the presence of a suitable catalyst to form nitric oxide and water vapor: 4 NH 3 (g) + 5 O 2 (g) 4 NO(g) + 6 H 2 O(g) How many liters of NH 3 (g) at 850 C and 5.00 atm are required to react with 1.00 mol of O 2 (g) in this reaction? Answer: 14.8 L

19 Sample Exercise Applying Dalton s Law of Partial Pressures A mixture of 6.00 g O 2 (g) and 9.00 g CH 4 (g) is placed in a 15.0-L vessel at 0 C. What is the partial pressure of each gas, and what is the total pressure in the vessel? Solution Analyze We need to calculate the pressure for two gases in the same volume and at the same temperature. Plan Because each gas behaves independently, we can use the ideal gas equation to calculate the pressure each would exert if the other were not present. The total pressure is the sum of these two partial pressures. Solve We first convert the mass of each gas to moles: We use the ideal-gas equation to calculate the partial Pressure of each gas: According to Dalton s law of partial pressures (Equation 10.12), the total pressure in the vessel is the sum of the partial pressures: Pt = P + P O = atm atm = atm 2 CH4

20 Sample Exercise Applying Dalton s Law of Partial Pressures Check A pressure of roughly 1 atm seems right for a mixture of about 0.2 mol O 2 and a bit more than 0.5 mol CH 4, together in a 15-L volume, because 1 mol of an ideal gas at 1 atm pressure and 0 C occupies about 22 L. Practice Exercise What is the total pressure exerted by a mixture of 2.00 g of H 2 (g) and 8.00 g of N 2 (g) at 273 K in a 10.0-L vessel? Answer: 2.86 atm

21 Sample Exercise Relating Mole Fractions and Partial Pressures A study of the effects of certain gases on plant growth requires a synthetic atmosphere composed of 1.5 mol percent CO 2, 18.0 mol percent O 2, and 80.5 mol percent Ar. (a) Calculate the partial pressure of O 2 Solution Analyze For (a) we need to calculate the partial pressure of O 2 given its mole percent and the total pressure of the mixture. For (b) we need to calculate the number of moles of O 2 in the mixture given its volume (121 L), temperature (745 torr), and partial pressure from part (a). Solve (a) The mole percent is the mole fraction times 100. Therefore, the mole fraction of O 2 is Equation gives in the mixture if the total pressure of the atmosphere is to be 745 torr. (b) If this atmosphere is to be held in a 121-L space at 295 K, how many moles of O 2 are needed? Plan We calculate the partial pressures using Equation 10.16, and then use P O2, V, and T in the ideal-gas equation to calculate the number of moles of O 2. P O2 = (0.180)(745 torr) = 134 torr

22 Sample Exercise Relating Mole Fractions and Partial Pressures (b) Tabulating the given variables and converting To appropriate units, we have Solving the ideal-gas equation for n O2, we have Check The units check out, and the answer seems to be the right order of magnitude. Practice Exercise From data gathered by Voyager 1, scientists have estimated the composition of the atmosphere of Titan, Saturn s largest moon. The pressure on the surface of Titan is 1220 torr. The atmosphere consists of 82 mol percent N 2, 12 mol percent Ar, and 6.0 mol percent CH 4. Calculate the partial pressure of each gas. Answer: torr N 2, torr Ar, and 73 torr CH 4

23 Sample Exercise Calculating the Amount of Gas Collected over Water When a sample of KClO 3 is partially decomposed in the setup shown in Figure 10.15, the volume of gas collected is L at 26 C and 765 torr total pressure. (a) How many moles of O 2 are collected? (b) How many grams of KClO 3 were decomposed? Solution (a) Analyze We need to calculate the number of moles of O 2 gas in a container that also contains water vapor. Plan We are given values for V and T. To use the ideal-gas equation to calculate the unknown, n O2, we must know the partial pressure of O 2 in the system. We can calculate this partial pressure from the total pressure (765 torr) and the vapor pressure of water. Solve The partial pressure of the O 2 gas is the difference between the total pressure and the pressure of the water vapor at 26 C, 25 torr (Appendix B): P O2 = 765 torr 25 torr = 740 torr We use the ideal-gas equation to calculate the number of moles of O 2 :

24 Sample Exercise Calculating the Amount of Gas Collected over Water (b) Analyze We need to calculate the number of moles of reactant KClO 3 decomposed. Plan We can use the number of moles of O 2 formed and the balanced chemical equation to determine the number of moles of KClO 3 decomposed, which we can then convert to grams of KClO 3. Solve From Equation 10.17, we have 2 mol. The molar mass of KClO 3 is g/mol. Thus, we can convert the number of moles of O 2 from part (a) to moles of KClO 3 and then to grams of KClO 3 : Check The units cancel appropriately in the calculations. The numbers of moles of O 2 and KClO 3 seem reasonable, given the small volume of gas collected.

25 Sample Exercise Calculating the Amount of Gas Collected over Water Practice Exercise Ammonium nitrite, NH 4 NO 2, decomposes on heating to form N 2 gas: NH 4 NO 2 (s) N 2 (g) + 2 H 2 O(l) When a sample of NH 4 NO 2 is decomposed in the apparatus of Figure 10.15, 511 ml of N 2 gas is collected over water at 26 C and 745 torr total pressure. How many grams of NH 4 NO 2 were decomposed? Answer: 1.26 g

26 Sample Exercise Applying the Kinetic-Molecular Theory A sample of O 2 gas initially at STP is compressed to a smaller volume at constant temperature. What effect does this change have on (a) the average kinetic energy of the molecules, (b) their average speed, (c) the number of collisions they make with the container walls per unit time, (d) the number of collisions they make with a unit area of container wall per unit time? Solution Analyze We need to apply the concepts of the kinetic-molecular theory of gases to a gas compressed at constant temperature. Plan We will determine how each of the quantities in (a) (d) is affected by the change in volume at constant temperature. Solve (a) Because the average kinetic energy of the O 2 molecules is determined only by temperature, this energy is unchanged by the compression. (b) Because the average kinetic energy of the molecules does not change, their average speed remains constant. (c) The number of collisions with the walls per unit time increases because the molecules are moving in a smaller volume but with the same average speed as before. Under these conditions they must encounter a wall more frequently. (d) The number of collisions with a unit area of wall per unit time increases because the total number of collisions with the walls per unit time increases and the area of the walls decreases. Check In a conceptual exercise of this kind, there is no numerical answer to check. All we can check in such cases is our reasoning in the course of solving the problem.

27 Sample Exercise Applying the Kinetic-Molecular Theory Practice Exercise How is the rms speed of N 2 molecules in a gas sample changed by (a) an increase in temperature, (b) an increase in volume, (c) mixing with a sample of Ar at the same temperature? Answers: (a) increases, (b) no effect, (c) no effect

28 Sample Exercise Calculating a Root-Mean-Square Speed Calculate the rms speed of the molecules in a sample of N 2 gas at 25 C. Solution Analyze We are given the identity of a gas and the temperature, the two quantities we need to calculate the rms speed. Plan We calculate the rms speed using Equation Solve We must convert each quantity in our equation to SI units. We will also use R in units of J/mol-K (Table 10.2) to make the units cancel correctly. Comment This corresponds to a speed of 1150 mi/hr. Because the average molecular weight of air molecules is slightly greater than that of N 2, the rms speed of air molecules is a little lower than that for N 2.

29 Sample Exercise Calculating a Root-Mean-Square Speed Practice Exercise What is the rms speed of an atom in a sample of He gas at 25 C? Answer: m/s

30 Sample Exercise Applying Graham s Law An unknown gas composed of homonuclear diatomic molecules effuses at a rate that is times the rate at which O 2 gas effuses at the same temperature. Calculate the molar mass of the unknown and identify it. Solution Analyze We are given the rate of effusion of an unknown gas relative to that of O 2 and asked to find the molar mass and identity of the unknown. Thus, we need to connect relative rates of effusion to relative molar masses. Plan We use Equation 10.24, to determine the molar mass of the unknown gas. If we let r x and Μ x represent the rate of effusion and molar mass of the gas, we can write Solve From the information given, Thus, r x = r O2

31 Sample Exercise Applying Graham s Law Because we are told that the unknown gas is composed of homonuclear diatomic molecules, it must be an element. The molar mass must represent twice the atomic weight of the atoms in the unknown gas. We conclude that the unknown gas is I 2. Practice Exercise Calculate the ratio of the effusion rates of N 2 gas and O 2 gas. Answer: r N2 /r O2 = 1.07

32 Sample Exercise Using the van der Waals Equation If mol of an ideal gas were confined to L at 0.0 C, it would exert a pressure of atm. Use the van der Waals equation and Table 10.3 to estimate the pressure exerted by mol of Cl 2 (g) in L at 0.0 C. Solution Analyze We need to determine a pressure. Because we will use the van der Waals equation, we must identify the appropriate values for the constants that appear there. Plan Solving Equation for P, we have Solve Substituting n = mol, R = L-atm/mol-K, T = K, V = L, a = 6.49 L 2 -atm/mol 2, and b = L/mol:

33 Sample Exercise Using the van der Waals Equation Check We expect a pressure not far from atm, which would be the value for an ideal gas, so our answer seems very reasonable. Comment Notice that the term atm is the pressure corrected for molecular volume. This value is higher than the ideal value, atm, because the volume in which the molecules are free to move is smaller than the container volume, L. Thus, the molecules collide more frequently with the container walls. The term atm corrects for intermolecular forces. The intermolecular attractions between molecules reduce the pressure to atm.we conclude, therefore, that the intermolecular attractions are the main cause of the slight deviation of Cl 2 (g) from ideal behavior under the stated experimental conditions. Practice Exercise A sample of mol of CO 2 (g) is confined to a L container at C. Calculate the pressure of the gas using (a) the ideal-gas equation and (b) the van der Waals equation. Answers: (a) 7.47 atm, (b) 7.18 atm

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