Kelter P., Mosher M., Scott A. Chemistry. The Practical Science

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Ozone

Oxygen gas

Oxygen atom

The reaction of ozone with

UV-B light reduces the intensity

of UV-B at the surface of the Earth.

428 Chapter 10 The Behavior and Applications of Gases

Pressure

(torr)

Temperature

(K)

Major

Constituents

Minor

Constituents

125

500

Exosphere

, O

, NO

, O

, e

–

Thermosphere

180

280

218

1 × 10

–3

150

760

Mesosphere

Stratosphere

Troposphere

100

Altitude (km)

FIGURE 10.24

Two of the four layers of the

atmosphere contain ozone, O

band speciﬁes radiation with a wavelength of 400 nm to 320 nm; UV-B indicates

wavelengths of 320 nm to 290 nm; and UV-C, the most energetic, is the classiﬁ-

cation of wavelengths from 290 nm to 100 nm. The destruction of ozone occurs

photodissociation—that is, UV light of sufﬁcient energy (UV-B) causes the

molecule to dissociate, forming oxygen gas and atomic oxygen:

(g)

−→

(g) + O(g)

There is also a mechanism for the formation of ozone, which also contributes to

the ﬁltering of UV radiation from the sun. Shorter-wavelength UV radiation

(UV-C) provides sufﬁcient energy for the O

to separate into oxygen atoms:

(g)

−→

2O(g)

These highly reactive individual atoms then recombine with O

to form ozone

along with the release of heat

(g) + O(g) n O

(g)

The resulting combination of these three equations gives rise to the Chapman

mechanism, a useful description of the chemistry involving oxygen and ozone in

the middle stratosphere.

The natural balance between ozone photodissociation and ozone formation

keeps most harmful UV radiation away from the Earth’s surface. Within the last

30 years, however, activities related to industrialization have led to a signiﬁcant

reduction in the stratospheric ozone layer over many parts of the planet, as

shown in Figure 10.25. Most notable is the discovery of a “hole”in the ozone layer

over Antarctica.

Sherwood Rowland and Mario Molina suggested in 1974 that

chloroﬂuorocar-

bons (CFCs)

, used in refrigeration and in the formation of polymer foams, were

the culprits in the destruction of the ozone layer. Table 10.7 lists common CFCs

and related compounds, called

halons, along with their uses.

CFCs were initially much sought after as refrigerants. This class of com-

pounds is chemically inert, nonﬂammable, inexpensive, and quite stable at the

Earth’s surface. These properties make them very suitable for use in homes and

automobiles. However, as they are released into the atmosphere, they travel up-

ward to the stratosphere without reacting with other compounds. Then, in the

stratosphere, they are bombarded with high-energy UV radiation and do react.

This produces highly reactive chlorine atoms that react with ozone as well as with

individual oxygen atoms that come from ozone.

Cl(g) + O

(g) n ClO(g) + O

(g)

ClO(g) + O(g) n Cl(g) + O

(g) .

overall reaction: O

(g) + O(g) n 2O

(g)

10.9 Industrialization: A Wonderful, Yet Cautionary, Tale 429

FIGURE 10.25

Total Ozone Mapping Spectrometer

(TOMS) data from satellites orbiting the

Earth give us a picture of the damage

done by ozone depleting substances.

These images from the period 1979 to

1999 show the development of a large

area of decreased ozone in the strato-

sphere over Antarctica.

UV-C + O

+ O

V-B

The Chapman mechanism.

Common Chloroﬂuorocarbons and Halons

Formula Common Symbol Major Uses

CCl

F CFC-11 Polymer foams, refrigeration, air conditioning

CCl

CFC-12 Polymer foams, refrigeration, air conditioning,

aerosols, food–freezing solvents

CCl

FCClF

CFC-113 Solvent

CBrClF

halon-1211 Portable ﬁre extinguishers

halon-2402 Fire extinguishers

TABLE 10.7

430 Chapter 10 The Behavior and Applications of Gases

As we monitor the decrease in stratospheric ozone with

growing concern, we also note the localized increase in

surface ozone, because this gas is an irritant to the eyes,

nose, and throat and damages cells in plants. Where does

tropospheric ozone come from? We typically note its

presence in the metropolitan areas of industrialized na-

tions. More speciﬁcally, the reaction mechanism for the

accumulation of ozone at the surface involves the release

of nitric oxide from automobiles without catalytic con-

verters. This is due to the unwanted combustion of ni-

trogen gas (present in the air) inside the combustion

chambers of automobiles:

(g) + O

(g) n 2NO(g)

Unfortunately, even automobiles with catalytic convert-

ers produce small amounts of NO. The released NO

reacts with oxygen to make nitrogen dioxide:

2NO(g) + O

(g) n 2NO

(g)

Then, when the Sun shines on a city ﬁlled with NO

the molecule decomposes to produce highly reactive

oxygen atoms:

sunlight

(g)

−−−−−→

NO(g) + O(g)(reaction A)

The atomic oxygen combines with oxygen gas in the

same way that it does in the stratosphere to produce

ozone:

(g) + O(g) n O

(g) (reaction B)

Then ozone is consumed in this reaction:

NO(g) + O

(g) n NO

(g) + O

(g) (reaction C)

The resulting combination of reactions A, B, and C that

produce and consume ozone at ground level is known

as a null cycle. No net reaction occurs in a null cycle.

Instead, the reactions are governed speciﬁcally by the

relative concentrations of the individual reactants and

products. Especially important is the presence of or-

ganic (carbon-based) pollutants and heat, such as are

present on a hot summer day. These pollutants create

highly reactive organic compounds that react faster with

NO than does O

. In the presence of these organic pol-

lutants, the null cycle is modiﬁed and excess O

remains!

What is the effect of this “nanoworld” process on the

“macroworld”? This is the impact of heat and pollution

in urban areas. Ozone accumulates at the Earth’s sur-

face, especially in large cities that have lots of motor

vehicle trafﬁc. Even smaller metropolitan areas, such as

the “Triad” cities in North Carolina shown in Figure

10.26, have experienced an increase in ozone concentra-

tions due to automobile exhaust.

NanoWorld / MacroWorld

The big effects of the very small:

Accumulation of ozone at the earth’s surface

Smog is a common annoyance—and often a hazard—in

industrialized cities. Its production is governed by a null

cycle that produces ozone in the presence of sunlight.

Air quality index (AQI)

Apr. 1

100

125

150

175

200

225

250

Apr. 15

Apr. 29

May 13

May 27

June 10

June 24

July 8

July 22

Aug. 5

Aug. 19

Sept. 2

Sept. 16

Sept. 30

Oct. 14

Oct. 28

2002 Date

Very unhealthful

Unhealthful

Unhealthful for sensitive groups

Moderate

Good

FIGURE 10.26

The ozone level in the “Triad” of cen-

tral North Carolina from April to Octo-

ber 2002. Even in this midsized region

(including the cities of Greensboro,

High Point, and Winston-Salem), air

quality is a signiﬁcant concern.

10.9 Industrialization: A Wonderful, Yet Cautionary, Tale 431

The chlorine atom is a catalyst in the destruction of ozone. Each chlorine atom

can catalyze the breakdown of perhaps hundreds of thousands of ozone

molecules.

Rowland and Molina shared the 1995 Nobel Prize in chemistry for their

groundbreaking research and predictions about the fate of the ozone layer. Since

that time, substitute compounds called HCFCs have been developed, including

CHF

Cl and CF

CFH

. These react in the lower atmosphere to release HCl and

HF, never making it to the stratosphere. The loss of ozone in the stratosphere,

combined with the generation of unhealthful ozone levels in large cities, is a con-

tinued cause for concern and a focus of research. The very good news is that as a

result of worldwide action to reduce CFC emissions, we may have turned the cor-

ner. In 2005, scientists found a slight worldwide increase in the amount of ozone

in the stratosphere. Many decades of continued diligence will be required before

the damage to the ozone layer can be completely reversed, yet we have made an

important start.

The Greenhouse Effect

The greenhouse effect is caused by the accumulation, in the atmosphere, of gases

that permit light to enter but prevent some energy as heat from exiting, much like

a plant greenhouse. Methane and CFCs are examples of greenhouse gases, al-

though they are not present in sufﬁcient concentration to be of real concern at

this time. A potent new greenhouse gas named triﬂuoromethyl sulfur pentaﬂuo-

ride (SF

) has recently been identiﬁed.

Carbon dioxide is entering the atmosphere in much greater amounts than are

being used up, and this makes CO

the most important greenhouse gas.

The

carbon cycle, discussed in Chapter 12, has historically kept the carbon in

the atmosphere, seas, and land in balance over the long term. Human razing of

forests has radically curtailed the number of plants that use up CO

photosynthesis:

6CO

(g) + 6H

O(g) n C

(s) + 6O

(g)

Burning carbon-based fuels, especially coal, oil, and natural gas, has added CO

to the environment, and the annual release of carbon into the atmosphere con-

tinues to rise (see Figure 10.27). There is general agreement among scientists

The greenhouse effect keeps the Earth warm. The gases in the atmosphere absorb infrared energy emitted

from the Earth and reemit it to the Earth, much as the panes of glass keep heat in a greenhouse.

Application

HEMICAL ENCOUNTERS:

The Greenhouse

Effect

432 Chapter 10 The Behavior and Applications of Gases

that the 100-year increase in the Earth’s average surface temperature, shown in

Figure 10.28, is at least partially caused by the greenhouse effect.

It is often said that the problems caused by the use of chemistry can also be

ﬁxed by the use of chemistry. Some good signs are on the horizon. As shown in

Figure 10.29, the fraction of global energy consumption derived from the burn-

ing of coal and wood is declining.

The interaction among gases in the atmosphere is so complex that we cannot

conﬁdently predict the long-term atmospheric effects of industrialization. The

best models indicate that as we produce gases that meet the needs and the wants

of the nearly 7 billion people on Earth, we must be mindful of the possible impact

of our current consumption. We are, after all, the keepers of the global commons.

800 1000 1200

Ice-core data

Atmospheric measurements

1400 1600 1800 2000

260

280

300

320

340

360

380

Carbon dioxide concentration (parts per million)

FIGURE 10.27

The amount of carbon released as CO

into the atmosphere continues to rise

along with our seemingly insatiable

appetite for fossil fuels.

GREENHOUSE GASES

Atmospheric carbon dioxide in parts per million

1880: 290.7

1999: 368.4

200

250

300

350

400

1900 1920 1940 1960 1980 19991880

Data prior to 1959 is derived from ice core samples.

SURFACE TEMPERATURES

Expressed as departures from the 1880–1998 average

1880 1900 1920 1940 1960 1980 1999

1.6

1.1

.54

-.54

Farenheit

-1.1

FIGURE 10.28

There is general, though not universal,

agreement among scientists that the

100-year increase in the Earth’s average

surface temperature, shown here, is at

least partially caused by the greenhouse

effect.

100

1855

Source: International Institute for Applied Systems Analysis.

1875 1895 1915 1935 1955 1975 1995

Ye a r

Wood

Share of global

energy consumption (%)

Coal

Oil

Natural gas

Hydropower

Nuclear

FIGURE 10.29

Advances in chemistry have changed

the way in which energy is produced

and consumed.

Key Words 433

The Bottom Line

■

Industrial gases are used in manufacturing, medi-

cine, and other industries. (Section 10.1)

■

Gases present at low density all behave in about the

same way. This permits us to generalize about their

behavior in a manner that is not possible with solids

and liquids. (Section 10.1)

■

Gases that behave as though each particle has no

interactions with any other are called ideal gases.

(Section 10.1)

■

High pressure, low temperature, and intermolecular

forces of attraction between gas molecules cause

gases to deviate from ideal behavior. (Section 10.1)

■

Real gases deviate from ideal behavior, especially at

low temperatures and high pressures. (Section 10.1)

■

Pressure is a measure of force per unit area. Pres-

sure is the result of collisions of gas molecules with

the walls of a container. (Section 10.2)

■

Several units of pressure exist, including the pascal

(the SI unit), atmosphere, mm Hg, in Hg, torr, bar,

psi, and psig. (Section 10.2)

■

In current usage, the standard atmosphere (1 atm)

is different from standard pressure (1 bar, or about

0.987 atm). (Section 10.2)

■

Dalton’s law of partial pressures is concerned with

the contribution of each gas to the pressure of the

entire gas mixture. (Section 10.3)

■

Avogadro’s law deals with the relationship between

the volume and number of moles of ideal gases, at

constant temperature and pressure. (Section 10.4)

■

Boyle’s law expresses the relationship between vol-

ume and pressure of an ideal gas, at constant num-

ber of moles and temperature. (Section 10.4)

■

Charles’s law expresses the relationship between the

volume and temperature of an ideal gas, at constant

number of moles and pressure. (Section 10.4)

■

The ideal gas equation combines the gas laws to

interrelate the pressure, volume, amount, and tem-

perature of an ideal gas. (Section 10.5)

■

The ideal gas equation can be applied to ﬁnd the

molar mass and the density of an ideal gas.

(Section 10.5)

■

Several scientists have made mathematical models

to account for nonideal behavior. J. D. van der

Waals’s model is the most commonly used because

it is relatively simple and takes into account correc-

tions for pressure and volume. (Section 10.5)

■

The molar mass and the density of an ideal gas are

directly proportional. (Section 10.6)

■

The kinetic-molecular theory shows how it is possi-

ble to start from some elementary constructs about

gas behavior and derive the gas laws, therefore

showing the consistency between theory and exper-

iment. (Section 10.7)

■

Graham’s law of effusion shows the inverse rela-

tionship between the speed of a gas and its molar

mass. (Section 10.8)

■

The effects of industrialization on the atmosphere

are currently being debated. Ozone levels and the

greenhouse effect are two issues of greatest social

concern. (Section 10.9)

Key Words

absolute zero The temperature obtained by extrapola-

tion of a plot of gas volume versus temperature to

the “zero volume” point. The lowest temperature

possible, which is 0 K, or –273.15°C. (p. 409)

atmosphere (atm) A unit used to measure pressure.

1 atm = 760 mm Hg. (p. 399)

average speed The sum of the molecular speeds of

each molecule divided by the number of molecules.

(p. 424)

Avogadro’s law Equal amounts of gases occupy the

same volume at constant temperature and pressure.

(p. 404)

bar A unit used to measure pressure. 1.01325 bar =

1 atmosphere. (p. 399)

Boyle’s law The volume of a ﬁxed amount of gas is

inversely proportional to its pressure at constant

temperature. (p. 406)

carbon cycle The natural process that describes how

carbon atoms are moved among the land, sea, and

atmosphere. (p. 431)

Charles’s law The volume of a ﬁxed amount of gas is

directly proportional to its temperature in kelvins at

a constant pressure. (p. 408)

chloroﬂuorocarbons (CFCs) Compounds containing only

carbon, chlorine, and ﬂuorine. CFCs are typically

used as solvents and refrigerants and are known to

be harmful to the ozone layer. (p. 429)

combined gas equation The combination of Avogadro’s

law, Boyle’s law, and Charles’s law. (p. 411)

434 Chapter 10 The Behavior and Applications of Gases

Dalton’s law of partial pressures The total pressure of a

mixture of gases is the simple sum of the individual

pressures of all the gaseous components. (p. 401)

diffusion The process by which two or more substances

mix. (p. 426)

effusion The process by which a gas escapes through a

small hole. (p. 425)

electrolysis A process in which an electric current pass-

ing through a solution causes a chemical reaction.

(p. 397)

force The mass of an object multiplied by its accelera-

tion. (p. 397)

Graham’s law of effusion The rates of the effusion of

gases are inversely proportional to the square roots

of their molar masses. (p. 426)

greenhouse effect The re-radiation of energy as heat

from gases in the atmosphere back toward Earth. An

increase in the greenhouse effect leads to an increase

in global temperatures. (p. 431)

halons Compounds containing carbon and halogens.

Halons are typically used in ﬁre extinguishing appli-

cations but are considered harmful to the ozone

layer. (p. 429)

ideal gases Gases that behave as though each particle

has no interactions with any other particle and

occupies no molar volume. (p. 396)

ideal gas constant R, the constant used in the ideal gas

equation. R = 0.08206 L·atm/mol·K. (p. 414)

ideal gas equation The equation relating the pressure,

volume, moles, and temperature of an ideal gas.

(p. 414)

in Hg A unit used to measure pressure. 29.921 in Hg

equals 1 atmosphere. (p. 399)

industrial gases Gases used in industrial or commercial

applications. (p. 395)

kinetic-molecular theory A theory explaining the rela-

tionship between kinetic energy and gas behavior.

This theory makes possible derivation of the ideal

gas laws. (p. 422)

mm Hg A unit used to measure pressure. 760 mm Hg

represents the height of a column of mercury that

can be supported by 1 atm pressure. (p. 399)

modiﬁed atmosphere packaging (MAP) The replacement

of the air within food packaging with gases in order

to prolong the shelf life of the food. (p. 401)

most probable speed The most likely speed of a mole-

cule from among a group of molecules. (p. 424)

newton (N) The basic SI unit of force.

1 N = 1 kg·m·s

−2

.(p. 397)

partial pressure The pressure exerted by a single com-

ponent of a gaseous mixture. (p. 401)

pascal (Pa) The basic SI unit of pressure.

1 Pa = 1 N·m

−2

.(p. 398)

photodissociation A chemical reaction where light, of

sufﬁcient energy, causes the cleavage of a bond in a

molecule. (p. 428)

pounds per square inch (psi) A unit of pressure.

14.7 psi = 1 atm. (p. 399)

pounds per square inch gauge (psig) A unit of pressure

expressing the pressure of a gas in excess of the

standard atmospheric pressure. Also known as gauge

pressure. 14.7 psig = 2 atm. (p. 399)

pressure The force per unit area exerted by an object

on another. (p. 398)

root-mean-square (rms) speed The square root of the

sum of the squares of the individual speeds of parti-

cles divided by the number of those particles.

(p. 424)

standard atmosphere Exactly 1 atm, the pressure that

supports 760 mm of mercury. (p. 399)

standard pressure Deﬁned as 1 bar, although chemists

typically use 1 atm. (p. 399)

standard temperature For gases, deﬁned to be 0°C, or

273 K. (p. 400)

standard temperature and pressure (STP) 1 atm and 273 K.

(p. 400)

standard ambient temperature and pressure (SATP) 1 bar

and 298 K. (p. 400)

torr A unit used to measure the pressure of gases.

1torr = 1 mm Hg. (p. 399)

van der Waals equation The equation that corrects the

gas laws for gases that deviate from ideal behavior.

(p. 417)

Focus Your Learning

The answers to the odd-numbered problems and some selected

problems appear at the back of the book, as represented by the

blue numbering.

10.1 The Nature of Gases

Skill Review

1. Deﬁne the intermolecular forces typically experienced by

molecules in the gaseous state.

2. Explain why we never see signs on delivery trucks that read

“Danger—Compressed Liquid.”

3. What conditions tend to favor ideal behavior in gases?

4. What are some characteristics common in nonideal gases?

10.2 Production of Hydrogen and the Meaning

of Pressure

Skill Review

5. If your chemistry textbook had a mass of 0.89 kg, how much

force, in newtons, would the textbook exert on a desk in your

classroom?

6. If your textbook (mass = 0.89 kg) measured 27 cm × 23 cm,

what pressure, in pascals, would the book exert on the desk?

7. How much force, in newtons, would an 86-kg man exert on

the ground?

8. A 115-lb ﬂight executive wears spike heels every day to work.

If each heel has an area of 4.0 cm

, what is the pressure

exerted by each heel in pascals? (Assume all of the woman’s

weight is placed on the heels of her shoes.)

9. The pressure of a gas is measured to be 797 mm Hg. Convert

this pressure into units of Pa, kPa, atm, torr, bar, and in Hg.

10. The pressure of a gas is measured to be 0.750 atm. Convert

this pressure into units of Pa, kPa, torr, bar, and mm Hg.

11. An astronaut weighs 145 lb on the planet Earth, where the

gravitational pull is 9.8 m/s

. She is sent on a mission to a

planet whose gravitational pull is only 0.33 m/s

. What would

a scale on this planet indicate that the astronaut weighs?

12. What is the mass in kg and weight in lb of a person who

exerts a force of 375 N on a scale?

13. The air in a car tire exerts a pressure of 22.3 psig. What is

the force exerted by the air if the total area inside the tire is

567 cm

14. A full book bag that exerts a force of 12 N on a table takes up

an area that measures 25 cm × 36 cm. Calculate the amount

of pressure exerted by the bag.

Chemical Applications and Practice

15. Small inﬂatable sleeping mattresses make camping enjoyable

for many outdoor enthusiasts. If the air pressure of a com-

fortable mat is measured as 17.5 psig, what is the actual

pressure that the air is exerting on the inside walls of the mat?

What is the pressure if reported in units of atmospheres?

16. As the altitude increases, the air pressure decreases. To prove

this fact, a mountain climbing group measured the air pres-

sure on Mount Everest at 0.67 atm. What would this pressure

be if it were reported in units of kPa, bar, mm Hg, and psi?

10.3 Mixtures of Gases—Dalton’s Law and Food Packaging

Skill Review

17. The partial pressure of O

in a sample of air on a mountain-

top is 115 mm Hg. The oxygen makes up 21% of the gas

in the atmosphere. What is the total atmospheric pressure on

the mountaintop?

18. A gaseous mixture contains 54% N

, 39% O

, and 7% CO

volume. If the total pressure is 813 mm Hg at STP, what is the

partial pressure of each of the gases?

19. A total of 12.4 g of N

and 12.4 g of O

are combined in a

mixture that exerts a pressure of 1.23 atm. Calculate the

partial pressure of N

and that of O

in the mixture.

20. A mixture of gases contains CO

and O

in a 3-to-1 ratio. If

the partial pressure of CO

in the mixture is 2.34 atm, what is

the partial pressure of O

? What is the total pressure of the

mixture?

Chemical Applications and Practices

21. A common way to generate small amounts of oxygen in the

lab is by heating potassium chlorate in the presence of a

catalyst and collecting the oxygen by bubbling it through

water. If the total pressure of the collected gas was 785 mm Hg

at 27°C, what is the partial pressure of the collected oxygen?

The pressure of gaseous water at 27°C is 27 torr.

22. Producing hydrogen gas is becoming a very important tech-

nical and economic concern for food processing, and poten-

tially as a fuel for future transportation needs. If a chemical

engineer were developing a model for this purpose and col-

lected 4.25 g of H

over water at 25°C that had a total pressure

of 1.15 atm, what would be the partial pressure of the H

collected? The pressure of gaseous water at 25°C is 23.8 torr.

23. In a helium–neon laser, a mixture of helium and neon gas is

contained in a small tube used to generate coherent light. In

some applications, this type of laser is used to scan price codes

at the grocery store. If the internal gas mixture is 8.95% Ne

and 91.15% He by volume and the total pressure is main-

tained at 3.42 mm Hg, what are the partial pressures of the gas

components?

24. A balloon is ﬁlled with H

and O

for a demonstration. The

gases are added such that they will react completely when a

ﬂame is brought in contact with the balloon. Write the equa-

tion describing this reaction. If the total pressure of the bal-

loon is 3.4 atm, what would the partial pressure of each gas

need to be in order for the reaction to go to completion?

10.4 The Gas Laws—Relating the Behavior of Gases

to Key Properties

Skill Review

25. At STP, a 2.5-L sample of gas contains 4.5 mol of Ne. How

many moles would the sample contain if the volume in-

creased to 5.0 L? . . . if the volume decreased to 1.0 L?

26. A balloon containing 0.50 mol of Heis found to have a volume

of 1.75 L.Whatvolumewould the balloon haveif anadditional

0.15 mol of He were added? What volume would the balloon

haveif 0.23 molof Hewereremovedfromthe balloon?Assume

that the temperature and pressure remain constant.

27. A container holds 4.70 L of air at a pressure of 861 torr. If the

gas in the container were lowered to standard pressure, what

would be its volume? If the pressure of the gas were reduced

to 400 torr, what would be its volume? Assume the tempera-

ture remains constant.

28. A balloon holds 2.50 L of air at a pressure of 19.5 psi. If the

balloon were squeezed to a volume of 1.00 L, what would be

the new pressure of the gas? If the balloon were stretched to a

volume of 5.00 L, what would be the new pressure of the gas?

Assume the temperature remains constant.

29. A balloon containing 1.00 L of CO

at 72°F is placed in a

freezer at −10°F. What is the new volume of the balloon? If

the balloon is placed in the oven at 250°F, what is the new vol-

ume of the balloon? Assume the pressure remains constant.

30. The volume of a ﬁxed quantity of gas at 25°C is changed from

0.750 L to 5.57 L. What is the resulting temperature of the gas

if the pressure remains constant? If the volume of the origi-

nal sample were reduced to 0.150 L, what would be the re-

sulting temperature? Assume the pressure remains constant.

31. A 0.47-mol sample of gas at 37°C occupies 3.20 L at 2839 mm

Hg. What volume would the same gas occupy at STP?

Focus Your Learning

435

Whipped cream in a bottle uses N

as the propellant. Does the electro-

static potential map for N

O tell us

anything about how to draw the

Lewis dot structure?

436 Chapter 10 The Behavior and Applications of Gases

32. Consider this experiment, wherein a gas is heated in a sealed

chamber of ﬁxed volume. What is the new pressure of the

gas?

33. A balloon containing 0.15 g of H

at STP is pressurized to

32.0 psi. What is the new temperature of the gas inside the

balloon? Assume the volume remains constant.

34. A CO gas cylinder possesses a volume of 45.0 L of gas at

2.20 × 10

psi. If the valve breaks and the gas is immediately

released from the cylinder, what is the new volume of the gas?

Assume the temperature remains constant and standard

pressure is attained.

Chemical Applications and Practice

35. Pressurized N

O gas can be used to provide the “inﬂating”

power for canned whipped cream. If such a can that is used to

provide the topping for a sundae contains 0.58 g of N

O and

corresponds to 0.217 L, how many liters of N

O will 0.33 g of

O ﬁll at the same temperature and pressure?

pressure is 0.988 atm, what pressure will be required to com-

press the gas mixture to 50.0 mL? Assume the temperature

remains constant.

37. The molecule nitrogen monoxide (NO) plays several vital

roles in animal physiology. Nitrogen monoxide helps regu-

late blood pressure, inﬂuences blood clotting, and also inﬂu-

ences the immune system. A researcher studying this impor-

tant gas isolates a small, 0.150-mL sample at 37°C. If this

sample is cooled to room temperature (25°C), what volume

will the sample have if the pressure remains constant?

38. Because ozone (O

) has beneﬁcial properties at high altitudes

and harmful properties at low altitudes, atmospheric scien-

tists often study it. If a 50.0-mL sample of ozone collected at

high altitude has a temperature of −25°C is brought back to

the lab to study (where the temperature is 25°C), what will

the new volume of the gas be? Assume the pressure remains

constant.

10.5 and 10.6 The Ideal Gas Equation and Its Applications

Skill Review

39. Determine the value of the ideal gas constant, R, to four sig-

niﬁcant ﬁgures if 1.00 mol of gas at 1.00 atm and 273 K oc-

cupies 22.4 L. Repeat the calculation, but use each of the fol-

lowing as the label for the equivalent amount of pressure.

a. Pa b. mm Hg c. psi

40. A researcher explores the ideal gas equation. He measures the

pressure of a system at different temperatures (keeping the

volume and moles of the gas constant) and creates a plot of

P versus T. What is the slope of the line that results if the vol-

ume is 1.0 L when 1.0 mol of gas is used in the exploration?

41. When inﬂated with 6.0 ×10

g of helium, a weather balloon

has a pressure of 0.955 atm at 22.6°C. What volume does the

balloon occupy?

42. What would be the pressure of H

in a tank that has a volume

of 25 L, if it contained 45 g of H

at 25°C?

43. What is the volume of 44.0 g of CO

at STP?...at 39°C and

0.500 atm?

44. What is the temperature of 85.0 g of N

if the volume is

7.49 L at 850 mm Hg?

45. How many moles of Ar would be found in a 4.33-L balloon

at STP?

46. A gas is trapped in a 1.00-L ﬂask at 27°C and has a pressure of

0.955 atm. The mass of the gas is measured at 1.95 g. Of

these, which is most likely to be the identity of the gas: NO,

, or N

47. An unknown gas has a density of 0.600 g/L at 743 mm Hg

and 66°C. What is the molar mass of the unknown gas?

48. What is the density of CO

at STP? ...ofN

at STP?

49. At STP, the density of methane (CH

) is 0.714 g/L. What is

the density of methane at 25°C and 1.15 atm?

50. If a 500.0-mL sample of air at 26.5°C and 698 mm Hg has a

mass of 0.543 g, what is the average molar mass of air?

51. A typical breath may have a volume of 450 mL. If the air you

breathed were 21% oxygen gas, how many molecules of

oxygen would you inhale at 37°C and 0.922 atm?

1 atm

273 K

5.44 mol of gas 5.44 mol of gas

? atm

500°C

36. Inside the cylinder of an automobile engine, a mixture of

gasoline vapor and oxygen is combusted by a spark. However,

the gaseous mixture is ﬁrst pressurized by the action of a

moving piston that decreases the volume of the gas. If the ini-

tial volume of the cylinder is approximately 485 mL and the

52. If you were able to inhale 2.50 L of air at STP, how many

moles of N

would you be breathing? Assume that air is com-

posed of 79% nitrogen and 21% oxygen.

53. A gas at 301 K and 0.97 atm has a density of 4.48 g/L. What is

the molar mass of the gas?

54. At 27.0°C and 802 torr, a gas has a molar mass of 62.0 g/mol.

What is the density of this gas? What would be the density of

the gas if the pressure were changed to 1.00 atm?

55. A 5.00-g sample of a gas has a volume of 85.0 mL at 20.0°C

and 1.00 atm. What is the molar mass of this gas? Would

the molar mass change if the temperature were increased to

23°C?

56. What is the density of CO

at 50°C and 0.44 atm?

Chemical Applications and Practice

57. A student working in the lab forgets to record the tempera-

ture at which she obtained 0.675 g of O

(g). She did, how-

ever, report the pressure as 745 mm Hg and the volume as

478 mL. What would have been the temperature at those

conditions?

58. Ozone (O

) in the stratosphere helps reduce the amount

of ultraviolet radiation reaching the surface of the Earth. As-

suming a temperature of −25°C and a partial pressure due

to ozone of 1.2 ×10

−7

atm, how many ozone molecules

would be present in 1.00 L of air in the stratosphere?

59. The metabolism of glucose is the main source of energy for

humans. This equation shows the overall reaction for the

process:

(s) + 6O

(g) n 6CO

(g) + 6H

O(g)

Calculate the volume of CO

produced at 37°C and 1.00 atm

when 10.0 g glucose is oxidized.

60. Small portable lighters use compressed butane (C

)as

theirfuel.Whenbutane iscombusted,this reactiontakesplace:

(g) + 13O

(g) n 8CO

(g) + 10H

O(g)

a. At 27.0°C and 0.888 atm, how many liters of oxygen are

required for every 1.00 g of butane to react completely?

b. If the system is compressed to 3.50 atm, keeping the

temperature at 27.0°C, how many liters of oxygen will be

required to react completely with 1.00 g of butane?

61. An industrial process can be used to change ethylene into

ethanol. The process uses a catalyst, high pressures (6.8 MPa),

and temperatures of 3.00 × 10

K. What is the density of

ethanol vapor under these conditions?

62. The familiar aroma of garlic comes mainly from the com-

pound diallyl disulﬁde. Based on these measurements, deter-

mine the molar mass of this pungent compound. At 298 K,

a 105-mL container holds 0.618 g of diallyl disulﬁde with a

pressure of 0.987 atm.

63. Argon gas is being used in some incandescent light bulbs to

extend the life of the tungsten ﬁlament. If a light bulb has a

volume of 200.0 mL and contains 0.100 g of Ar at 25°C, what

is the pressure inside the bulb?

64. An aerosol can contains carbon dioxide as its propellant. If

the can has an internal pressure of 1.3 atm at 25.0°C, what

Focus Your Learning 437

pressure would it have if the temperature were accidentally

raised to 450°C? (Assume that the can has not yet burst,

although this would probably be a very close call.)

65. The Haber process is used to produce ammonia (NH

) for

use in fertilizing corn ﬁelds in the Midwest. The reaction is

(g) + 3H

(g) n 2NH

(g)

a. Which, if either, is the limiting reagent when 2.00 L of N

at STP is prepared to react with 2.00 L of H

at 1120 mm Hg

and 21.0°C?

b. If 2.00 L of the gases, both at STP, reacted via the Haber

process, how many liters of ammonia would be produced?

66. Much of the electricity in the United States is produced from

the combustion of coal. However, some of the coal deposits

contain FeS

(iron pyrite) as a contaminant. During the

combustion of coal, this impurity also burns and produces

harmful sulfur dioxide gas (SO

). How many liters of SO

are

produced from the burning of 1.00 kg of FeS

at 758 torr

and 275°C?

FeS

(s) + 2O

(g) n Fe(s) + 2SO

(g)

67. A gaseous sulfur compound causes the pungent aroma of

rotten eggs. If 1.60 g of the gas were collected in a 1.00-L ves-

sel at 1065 torr and 89.0°C, what would be the molar mass of

the trapped gas?

68. One common way to obtain metal samples from their im-

pure oxide ores is to react the oxide with carbon. Generically,

the equation can be written this way:

2MO(s) + C(s) n 2M(s) + CO

(g)

If 5.00 g of an unknown metal oxide (MO) reacted with

excess carbon and formed 0.738 L of CO

at 200.0°C and

0.978 atm, what is the identity of the metal?

69. The halogen gases are caustic and toxic. Therefore, they must

be handled with extreme care. Which halogen gas has a den-

sity of approximately 1.70 g/L at STP?

70. Hydrogen has an isotope known as deuterium (D) that is

made up of a proton, a neutron, and an electron. Deuterium

is present in measurable amounts in all hydrogen-containing

substances. At STP, what are the densities of the three com-

pounds that could exist in a hydrogen gas sample: H

,HD,

and D

71. Neon lights actually do contain neon gas. But from where do

we get neon? Very small amounts of Ne are present in air, less

than 2.0 × 10

−3

% by volume. However, when air is liqueﬁed,

the sample can carefully be separated on the basis of boiling

points. This must be controlled, but because N

, and Ar

boil away in the range from 77 to 97 K, neon (with a boiling

point of 27 K) can be separated. If a manufacturer collected

1.00 lb of neon at 25°C and 789 torr, what volume would the

Ne occupy?

72. Assume you are pumping air into the tires of your mountain

bike before a ride. If the tire volume is 1.50 L at 28.5°C and

6.55 atm, how many moles of air have you put inside the tire?

How many molecules of air are in the tire? (Refer to Problem

50 for the average molar mass of air.)