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Unformatted text preview: Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website CHAPTER 17 Chemistry in the Atmosphere INTRODUCTION WE HAVE STUDIED BASIC DEFINITIONS IN CHEMISTRY, AND WE HAVE EX- 17.1 EARTH’S ATMOSPHERE AMINED THE PROPERTIES OF GASES, LIQUID, SOLIDS, AND SOLUTIONS. WE 17.2 PHENOMENA IN THE OUTER LAYERS OF THE ATMOSPHERE HAVE DISCUSSED CHEMICAL BONDING AND INTERMOLECULAR FORCES AND SEEN HOW CHEMICAL KINETICS AND CHEMICAL EQUILIBRIUM CON- 17.3 DEPLETION OF OZONE IN THE STRATOSPHERE CEPTS HELP US UNDERSTAND THE NATURE OF CHEMICAL REACTIONS. 17.4 VOLCANOES IT IS APPROPRIATE AT THIS STAGE TO APPLY OUR KNOWLEDGE TO THE 17.5 THE GREENHOUSE EFFECT STUDY OF ONE EXTREMELY IMPORTANT SYSTEM: THE ATMOSPHERE. 17.6 ACID RAIN ALTHOUGH EARTH’S 17.7 PHOTOCHEMICAL SMOG ATMOSPHERE IS FAIRLY SIMPLE IN COMPOSITION, ITS CHEMISTRY IS VERY COMPLEX AND NOT FULLY UNDERSTOOD. THE CHEM- 17.8 INDOOR POLLUTION ICAL PROCESSES THAT TAKE PLACE IN OUR ATMOSPHERE ARE INDUCED BY SOLAR RADIATION, BUT THEY ARE INEXTRICABLY CONNECTED TO NATURAL EVENTS AND HUMAN ACTIVITIES ON EARTH’S SURFACE. IN THIS CHAPTER WE WILL DISCUSS THE STRUCTURE AND COMPOSITION OF THE ATMOSPHERE, TOGETHER WITH SOME OF THE CHEMICAL PROCESSES THAT OCCUR THERE. IN ADDITION, WE WILL TAKE A LOOK AT THE MAJOR SOURCES OF AIR POLLUTION AND PROSPECTS FOR CONTROLLING THEM. 693 Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 694 CHEMISTRY IN THE ATMOSPHERE 17.1 TABLE 17.1 Composition of Dry Air at Sea Level GAS COMPOSITION (% BY VOLUME) N2 O2 Ar CO2 Ne He Kr Xe 78.03 20.99 0.94 0.033 0.0015 0.000524 0.00014 0.000006 EARTH’S ATMOSPHERE Earth is unique among the planets of our solar system in having an atmosphere that is chemically active and rich in oxygen. Mars, for example, has a much thinner atmosphere that is about 90 percent carbon dioxide. Jupiter, on the other hand, has no solid surface; it is made up of 90 percent hydrogen, 9 percent helium, and 1 percent other substances. It is generally believed that three billion or four billion years ago, Earth’s atmosphere consisted mainly of ammonia, methane, and water. There was little, if any, free oxygen present. Ultraviolet (UV) radiation from the sun probably penetrated the atmosphere, rendering the surface of Earth sterile. However, the same UV radiation may have triggered the chemical reactions (perhaps beneath the surface) that eventually led to life on Earth. Primitive organisms used energy from the sun to break down carbon dioxide (produced by volcanic activity) to obtain carbon, which they incorporated in their own cells. The major by-product of this process, called photosynthesis, is oxygen. Another important source of oxygen is the photodecomposition of water vapor by UV light. Over time, the more reactive gases such as ammonia and methane have largely disappeared, and today our atmosphere consists mainly of oxygen and nitrogen gases. Biological processes determine to a great extent the atmospheric concentrations of these gases, one of which is reactive (oxygen) and the other unreactive (nitrogen). Table 17.1 shows the composition of dry air at sea level. The total mass of the atmosphere is about 5.3 1018 kg. Water is excluded from this table because its concentration in air can vary drastically from location to location. Figure 17.1 shows the major processes involved in the cycle of nitrogen in nature. Molecular nitrogen, with its triple bond, is a very stable molecule. However, through biological and industrial nitrogen fixation, the conversion of molecular nitrogen into nitrogen compounds, atmospheric nitrogen gas is converted into nitrates and other compounds suitable for assimilation by algae and plants. Another important mechanism for producing nitrates from nitrogen gas is lightning. The steps are N2(g) 2NO(g) 2NO2(g) Lightning is a major contributor to nitrogen fixation. Back Forward electrical O2(g) 887777n 2NO(g) energy O2(g) 887777n 2NO2(g) H2O(l ) 887777n HNO2(aq) HNO3(aq) About 30 million tons of HNO3 are produced this way annually. Nitric acid is converted to nitrate salts in the soil. These nutrients are taken up by plants, which in turn are ingested by animals. Animals use the nutrients from plants to make proteins and other essential biomolecules. Denitrification reverses nitrogen fixation to complete the cycle. For example, certain anaerobic organisms decompose animal wastes as well as dead plants and animals to produce free molecular nitrogen from nitrates. The main processes of the global oxygen cycle are shown in Figure 17.2. This cycle is complicated by the fact that oxygen takes so many different chemical forms. Atmospheric oxygen is removed through respiration and various industrial processes (mostly combustion), which produce carbon dioxide. Photosynthesis is the major mechanism by which molecular oxygen is regenerated from carbon dioxide and water. Scientists divide the atmosphere into several different layers according to temperature variation and composition (Figure 17.3). As far as visible events are concerned, the most active region is the troposphere, the layer of the atmosphere which contains Main Menu TOC Study Guide TOC Textbook Website MHHE Website 17.1 695 EARTH’S ATMOSPHERE Atmospheric nitrogen Atmospheric fixation Fixed juvenile nitrogen Industrial fixation Protein Biological fixation Igneous rocks Nitrate reduction Plant and animal wastes, dead organisms Denitrification Nitrous oxide Ammonium Nitrite Nitrate To ground water FIGURE 17.1 The nitrogen cycle. Although the supply of nitrogen in the atmosphere is virtually inexhaustible, it must be combined with hydrogen or oxygen before it can be assimilated by plants, which in turn are consumed by animals. Juvenile nitrogen is nitrogen that has not previously participated in the nitrogen cycle. about 80 percent of the total mass of air and practically all of the atmosphere’s water vapor. The troposphere is the thinnest layer of the atmosphere (10 km), but it is where all the dramatic events of weather—rain, lightning, hurricanes—occur. Temperature decreases almost linearly with increasing altitude in this region. Above the troposphere is the stratosphere, which consists of nitrogen, oxygen, and ozone. In the stratosphere, the air temperature rises with altitude. This warming effect is the result of exothermic reactions triggered by UV radiation from the sun (to be discussed in Section 17.3). One of the products of this reaction sequence is ozone (O3), which, as we will see shortly, serves to prevent harmful UV rays from reaching Earth’s surface. In the mesosphere, which is above the stratosphere, the concentration of ozone and other gases is low, and the temperature decreases with increasing altitude. The thermosphere, or ionosphere, is the uppermost layer of the atmosphere. The rise in temperature in this region is the result of the bombardment of molecular oxygen and nitrogen and atomic species by energetic particles, such as electrons and protons, from the sun. Typical reactions are N2 88n 2N H° 941.4 kJ N 88n N Back Forward Main Menu TOC e H° 1400 kJ O2 88n O2 e H° 1176 kJ Study Guide TOC Textbook Website MHHE Website 696 CHEMISTRY IN THE ATMOSPHERE High-energy ultraviolet radiation O O2 H O2 H2O H2O O2 Ozone screen O2 CO2 O2 Volcanism Oxidative weathering 4FeO + O2 CO2 Photic zone H2CO3 – HCO3 + H+ 2CO2 CO O CO2 Phytoplankton H2O + CO2 O2 + 2CO OH O3 2Fe2O3 Sediments – 2HCO3 H2O CO2 – 3 Ca2+ CaCO3 Sediments FIGURE 17.2 The oxygen cycle. The cycle is complicated because oxygen appears in so many chemical forms and combinations, primarily as molecular oxygen, in water, and in organic and inorganic compounds. In reverse, these processes liberate the equivalent amount of energy, mostly as heat. Ionized particles are responsible for the reflection of radio waves back toward Earth. 17.2 PHENOMENA IN THE OUTER LAYERS OF THE ATMOSPHERE In this section we will discuss two dazzling phenomena that occur in the outer regions of the atmosphere. One is a natural event. The other is a curious by-product of human space travel. AURORA BOREALIS AND AURORA AUSTRALIS Violent eruptions on the surface of the sun, called solar flares, result in the ejection of myriad electrons and protons into space, where they disrupt radio transmission and provide us with spectacular celestial light shows known as auroras (Figure 17.4). These electrons and protons collide with the molecules and atoms in Earth’s upper atmosphere, causing them to become ionized and electronically excited. Eventually, the excited molecules and ions return to the ground state with the emission of light. For example, an excited oxygen atom emits photons at wavelengths of 558 nm (green) and between 630 nm and 636 nm (red): O* 88n O Back Forward Main Menu TOC Study Guide TOC h Textbook Website MHHE Website 17.2 FIGURE 17.3 Regions of Earth’s atmosphere. Notice the variation in temperature with altitude. Most of the phenomena shown here are discussed in the chapter. 697 PHENOMENA IN THE OUTER LAYERS OF THE ATMOSPHERE 500 km 1000°C 400 Space shuttle 300 Thermosphere 950°C 200 900°C Aurora borealis 700°C 100 165°C Shooting star –80°C Mesosphere 80 50 Troposphere Stratosphere 0°C Back Forward Main Menu Ozone layer –50°C Concorde 10 0 TOC Mt. Pinatubo 1°C Study Guide TOC Textbook Website MHHE Website 698 CHEMISTRY IN THE ATMOSPHERE FIGURE 17.4 Aurora borealis, commonly referred to as the northern lights. where the asterisk denotes an electronically excited species and h the emitted photon (see Section 7.2). Similarly, the blue and violet colors often observed in auroras result from the transition in the ionized nitrogen molecule: N2 * 88n N2 h The wavelengths for this transition fall between 391 and 470 nm. The incoming streams of solar protons and electrons are oriented by Earth’s magnetic field so that most auroral displays occur in doughnut-shaped zones about 2000 km in diameter centered on the North and South Poles. Aurora borealis is the name given to this phenomenon in the Northern Hemisphere. In the Southern Hemisphere, it is called aurora australis. Sometimes, the number of solar particles is so immense that auroras are also visible from other locations on Earth. EXAMPLE 17.1 The bond dissociation energy of O2 is 498.7 kJ/mol. Calculate the maximum wavelength (nm) of a photon that can cause the dissociation of an O2 molecule. Answer First we need to calculate the energy required to break up one O2 mole- cule: 498.7 103 J mol 1 mol 1023 molecules 6.022 The energy of the photon is given by E E h 8.281 10 h [Equation (7.2)]. Therefore 1015 s 1 Finally, we calculate the wavelength of the photon, given by Forward Main Menu TOC J molecule 8.281 10 19 J 6.63 10 34 J s 1.25 Back 19 Study Guide TOC c/ , as follows: Textbook Website MHHE Website 17.2 PHENOMENA IN THE OUTER LAYERS OF THE ATMOSPHERE 3.00 1.25 108 m/s 1015 s 1 2.40 10 699 7 m 240 nm In principle, any photon with a wavelength of 240 nm or shorter can dissociate an O2 molecule. Comment Similar problem: 17.11. PRACTICE EXERCISE Calculate the wavelength (in nm) of a photon needed to dissociate an O3 molecule: O3 88n O O2 H° 107.2 kJ THE MYSTERY GLOW OF SPACE SHUTTLES A human-made light show that baffled scientists for several years is produced by space shuttles orbiting Earth. In 1983, astronauts first noticed an eerie orange glow on the outside surface of their spacecraft at an altitude about 300 km above Earth (Figure 17.5). The light, which usually extends about 10 cm away from the protective silica heat tiles and other surface materials, is most pronounced on the parts of the shuttle facing its direction of travel. This fact led scientists to postulate that collision between oxygen atoms in the atmosphere and the fast-moving shuttle somehow produced the orange light. Spectroscopic measurements of the glow, as well as laboratory tests, strongly suggested that nitric oxide (NO) and nitrogen dioxide (NO2) also played a part. It is believed that oxygen atoms interact with nitric oxide adsorbed on (that is, bound to) the shuttle’s surface to form electronically excited nitrogen dioxide: O NO 88n NO2* As the NO2* leaves the shell of the spacecraft, it emits photons at a wavelength of 680 nm (orange). NO2* 88n NO2 h Support for this explanation came inadvertently in 1991, when astronauts aboard Discovery released various gases, including carbon dioxide, neon, xenon, and nitric oxide, from the cargo bay in the course of an unrelated experiment. Expelled one at a time, these gases scattered onto the surface of the shuttle’s tail. The nitric oxide caused the normal shuttle glow to intensify markedly, but the other gases had no effect on it. What is the source of the nitric oxide on the outside of the spacecraft? Scientists believe that some of it may come from the exhaust gases emitted by the shuttle’s rockFIGURE 17.5 The glowing tail section of the space shuttle viewed from inside the vehicle. Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 700 CHEMISTRY IN THE ATMOSPHERE ets and that some of it is present in the surrounding atmosphere. The shuttle glow does not harm the vehicle, but it does interfere with spectroscopic measurements on distant objects made from the spacecraft. 17.3 Photodissociation is the breaking of chemical bonds by radiant energy. DEPLETION OF OZONE IN THE STRATOSPHERE As mentioned earlier, ozone in the stratosphere prevents UV radiation emitted by the sun from reaching Earth’s surface. The formation of ozone in this region begins with the photodissociation of oxygen molecules by solar radiation at wavelengths below 240 nm: UV O2 88888n O 240 nm O (17.1) The highly reactive O atoms combine with oxygen molecules to form ozone as follows: O O2 M 88n O3 M (17.2) where M is some inert substance such as N2. The role of M in this exothermic reaction is to absorb some of the excess energy released and prevent the spontaneous decomposition of the O3 molecule. The energy that is not absorbed by M is given off as heat. (As the M molecules themselves become de-excited, they release more heat to the surroundings.) In addition, ozone itself absorbs UV light between 200 and 300 nm: UV O3 88n O A can of CFC used in an air conditioner. Back Forward O2 (17.3) The process continues when O and O2 recombine to form O3 as shown in Equation (17.2), further warming the stratosphere. If all the atmospheric ozone were compressed into a single layer at STP on Earth, that layer would be only about 3 mm thick! Although the concentration of ozone in the stratosphere is very low, it is sufficient to filter out (that is, absorb) solar radiation in the 200- to 300-nm range [see Equation (17.3)]. In the stratosphere, it acts as our protective shield against UV radiation, which can induce skin cancer, cause genetic mutations, and destroy crops and other forms of vegetation. The formation and destruction of ozone by natural processes is a dynamic equilibrium that maintains a constant concentration of ozone in the stratosphere. Since the mid-1970s scientists have been concerned about the harmful effects of certain chlorofluorocarbons (CFCs) on the ozone layer. The CFCs, which are generally known by the trade name Freons, were first synthesized in the 1930s. Some of the common ones are CFCl3 (Freon 11), CF2Cl2 (Freon 12), C2F3Cl3 (Freon 113), and C2F4Cl2 (Freon 114). Because these compounds are readily liquefied, relatively inert, nontoxic, noncombustible, and volatile, they have been used as coolants in refrigerators and air conditioners, in place of highly toxic liquid sulfur dioxide (SO2) and ammonia (NH3). Large quantities of CFCs are also used in the manufacture of disposable foam products such as cups and plates, as aerosol propellants in spray cans, and as solvents to clean newly soldered electronic circuit boards (Figure 17.6). In 1977, the peak year of production, nearly 1.5 106 tons of CFCs were produced in the United States. Most of the CFCs produced for commercial and industrial use are eventually discharged into the atmosphere. Main Menu TOC Study Guide TOC Textbook Website MHHE Website 17.3 FIGURE 17.6 Uses of CFCs. Since 1978, the use of aerosol propellants has been banned in the United States. DEPLETION OF OZONE IN THE STRATOSPHERE 701 Recycling feasible Recycling not feasible Foam insulation 20% Auto air conditioning 21% Commercial refrigeration 17% Other foam uses 13% 4% 11% Solvent cleaning 14% Aerosols Others (sterilization, household refrigeration) It can take years for CFCs to reach the stratosphere. Because of their reactive inertness, the CFCs slowly diffuse unchanged up to the stratosphere, where UV radiation of wavelengths between 175 nm and 220 nm causes them to decompose: CFCl3 88n CFCl2 ClO is called chlorine monoxide. Cl CF2Cl2 88n CF2Cl Cl The reactive chlorine atoms then undergo the following reactions: Cl O3 88n ClO ClO O 88n Cl O2 O2 (17.4) (17.5) The overall result [sum of Equations (17.4) and (17.5)] is the net removal of an O3 molecule from the stratosphere: O3 Cl is a homogeneous catalyst. Back Forward Main Menu O 88n 2O2 (17.6) The oxygen atoms in Equation (17.5) are supplied by the photochemical decomposition of molecular oxygen and ozone described earlier. Note that the Cl atom plays the role of a catalyst in the reaction mechanism scheme represented by Equations (17.4) and (17.5) because it is not used up and therefore can take part in many such reactions. One Cl atom can destroy up to 100,000 O3 molecules before it is removed by some other reaction. The ClO species is an intermediate because it is produced in the first elementary step [Equation (17.4)] and consumed in the second step [Equation (17.5)]. The above mechanism for the destruction of ozone has been supported by the detection of ClO in the stratosphere in recent years (Figure 17.7). As can be seen, the concentration of O3 decreases in regions that have high amounts of ClO. Another group of compounds that can destroy stratospheric ozone are the nitrogen oxides, generally denoted as NOx. (Examples of NOx are NO, NO2, N2O, and N2O5.) These compounds come from the exhausts of high-altitude supersonic aircraft and from human and natural activities on Earth. Solar radiation decomposes a substantial amount of the other nitrogen oxides to nitric oxide (NO), which participates in the destruction of ozone as follows: TOC Study Guide TOC Textbook Website MHHE Website CHEMISTRY IN THE ATMOSPHERE FIGURE 17.7 The variations in the concentrations of ClO and O3 with latitude. Chlorine monoxide (ppb by volume) 2.5 O3 1.0 2.0 1.5 0.5 1.0 ClO Ozone (ppm by volume) 702 0.5 0 0 63°S 72°S Latitude O3 88n O2 NO NO2 Overall: O3 88n NO2 O 88n NO O O2 O2 2O3 88n 3O2 In this case, NO is the catalyst and NO2 is the intermediate. Nitrogen dioxide also reacts with chlorine monoxide to form chlorine nitrate: ClO NO2 88n ClONO2 Chlorine nitrate is relatively stable and serves as a “chlorine reservoir,” which plays a role in the depletion of the stratospheric ozone over the North and South Poles. POLAR OZONE HOLES In the mid-1980s, evidence began to accumulate that an “Antarctic ozone hole” developed in late winter, depleting the stratospheric ozone over Antarctica by as much as 50 percent (Figure 17.8). In the stratosphere, a stream of air known as the “polar vortex” circles Antarctica in winter. Air trapped within this vortex becomes extremely cold during the polar night. This condition leads to the formation of ice particles known as polar stratospheric clouds (PSCs) (Figure 17.9). Acting as a heterogeneous catalyst, these PSCs provide a surface for reactions converting HCl (emitted from Earth) and chlorine nitrate to more reactive chlorine molecules: HCl ClONO2 88n Cl2 HNO3 By early spring, the sunlight splits molecular chlorine into chlorine atoms Cl2 h 88n 2Cl which then attack ozone as shown earlier. The situation is not as severe in the warmer Arctic region, where the vortex does not persist quite as long. Studies have shown that ozone levels in this region have de- Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 17.3 FIGURE 17.8 In recent years, scientists have found that the ozone layer in the stratosphere over the South Pole has become thinner. This map, based on data collected over a number of years, shows the depletion of ozone in purple. (Source: NASA/Goddard Space Flight Center). The trend continues after 1994. DEPLETION OF OZONE IN THE STRATOSPHERE 703 clined between 4 and 8 percent in the past decade. Volcanic eruptions, such as that of Mount Pinatubo in the Philippines in 1991, inject large quantities of dust-sized particles and sulfuric acid aerosols (see p. 494) into the atmosphere. These particles can perform the same catalytic function as the ice crystals at the South Pole. As a result, the Arctic hole is expected to grow larger during the next few years. Recognizing the serious implications of the loss of ozone in the stratosphere, nations throughout the world have acknowledged the need to drastically curtail or totally stop the production of CFCs. In 1978 the United States was one of the few countries to ban the use of CFCs in hair sprays and other aerosols. An international treaty—the Montreal protocol—was signed by most industrialized nations in 1987, setting targets for cutbacks in CFC production and the complete elimination of these substances by the year 2000. While some progress has been made in this respect, it is doubtful that poorer nations such as China and India can strictly abide by the treaty because of the FIGURE 17.9 Polar stratospheric clouds containing ice particles can catalyze the formation of Cl atoms and lead to the destruction of ozone. Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 704 CHEMISTRY IN THE ATMOSPHERE The OH radical is formed by a series of complex reactions in the troposphere that are driven by sunlight. importance of CFCs to their economies. Recycling could play a significant supplementary role in preventing CFCs already in appliances from escaping into the atmosphere. As Figure 17.6 shows, more than half of the CFCs in use are recoverable. An intense effort is under way to find CFC substitutes that are not harmful to the ozone layer. One of the promising candidates is called hydrochlorofluorocarbon-123, or HCFC-123 (CF3CHCl2). The presence of the hydrogen atom makes the compound more susceptible to oxidation in the lower atmosphere, so that it never reaches the stratosphere. Specifically, it is attacked by the hydroxyl radical in the troposphere: CF3CHCl2 OH 88n CF3CCl2 H2O The CF3CCl2 fragment reacts with oxygen, eventually decomposing to CO2, water, and hydrogen halides that are removed by rainwater. Unfortunately, the same hydrogen atom also makes the compound more active biologically than the CFCs. Laboratory tests have shown that HCFC-123 can cause tumors in rats, although its toxic effect on humans is not known. Another promising group of compounds that can substitute for CFCs are the hydrofluorocarbons (HFCs). Because they do not contain chlorine, HFCs will not promote the destruction of ozone even if they diffuse to the stratosphere. Examples of these compounds are: CF3CFH2, CF3CF2H, CF3CH3, and CF2HCH3. In particular, CF3CFH2 is already widely used in place of CFCs in air conditioning and refrigeration applications. Although it is not clear whether the CFCs already released to the atmosphere will eventually result in catastrophic damage to life on Earth, it is conceivable that the depletion of ozone can be slowed by reducing the availability of Cl atoms. Indeed, some chemists have suggested sending a fleet of planes to spray 50,000 tons of ethane (C2H6) or propane (C3H8) high over the South Pole in an attempt to heal the hole in the ozone layer. Being a reactive species, the chlorine atom would react with the hydrocarbons as follows: Cl C2H6 88n HCl C2H5 Cl C3H8 88n HCl C3H7 The products of these reactions would not affect the ozone concentration. A less realistic plan is to rejuvenate the ozone layer by producing large quantities of ozone and releasing it into the stratosphere from airplanes. Technically this solution is feasible, but it would be enormously costly and it would require the collaboration of many nations. Having discussed the chemistry in the outer regions of Earth’s atmosphere, we will focus in the next five sections on events closer to us, that is, in the troposphere. 17.4 VOLCANOES Volcanic eruptions, Earth’s most spectacular natural displays of energy, are instrumental in forming large parts of Earth’s crust. The upper mantle, immediately under the crust, is nearly molten. A slight increase in heat, such as that generated by the movement of one crustal plate under another, melts the rock. The molten rock, called magma, rises to the surface and generates some types of volcanic eruptions (Figure 17.10). An active volcano emits gases, liquids, and solids. The gases spewed into the atmosphere include primarily N2, CO2, HCl, HF, H2S, and water vapor. It is estimated that volcanoes are the source of about two-thirds of the sulfur in the air. On the slopes Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 17.5 THE GREENHOUSE EFFECT 705 FIGURE 17.10 A volcanic eruption on the island of Hawaii. of Mount St. Helens, which last erupted in 1980, deposits of elemental sulfur are visible near the eruption site. At high temperatures, the hydrogen sulfide gas given off by a volcano is oxidized by air: 2H2S(g) 3O2(g) 88n 2SO2(g) 2H2O(g) Some of the SO2 is reduced by more H2S from the volcano to elemental sulfur and water: 2H2S(g) SO2(g) 88n 3S(s) 2H2O(g) The rest of the SO2 is released into the atmosphere, where it reacts with water to form acid rain (see Section 17.6). The tremendous force of a volcanic eruption carries a sizable amount of gas into the stratosphere. There SO2 is oxidized to SO3, which is eventually converted to sulfuric acid aerosols in a series of complex mechanisms. In addition to destroying ozone in the stratosphere (see p. 700), these aerosols can also affect climate. Because the stratosphere is above the atmospheric weather patterns, the aerosol clouds often persist for more than a year. They absorb solar radiation and thereby cause a drop in temperature at Earth’s surface. However, this cooling effect is local rather than global, because it depends on the site and frequency of volcanic eruptions. 17.5 A dramatic illustration of the greenhouse effect is found on Venus, where the atmosphere is 97 percent CO2 and the atmospheric pressure is 9 106 Pa (equivalent to 89 atm). The surface temperature of Venus is about 730 K! Back Forward Main Menu THE GREENHOUSE EFFECT Although carbon dioxide is only a trace gas in Earth’s atmosphere, with a concentration of about 0.033 percent by volume (see Table 17.1), it plays a critical role in controlling our climate. The so-called greenhouse effect describes the trapping of heat near Earth’s surface by gases in the atmosphere, particularly carbon dioxide. The glass roof of a greenhouse transmits visible sunlight and absorbs some of the outgoing infrared (IR) radiation, thereby trapping the heat. Carbon dioxide acts somewhat like a glass roof, except that the temperature rise in the greenhouse is due mainly to the restricted air circulation inside. Calculations show that if the atmosphere did not contain carbon dioxide, Earth would be 30°C cooler! Figure 17.11 shows the carbon cycle in our global ecosystem. The transfer of carbon dioxide to and from the atmosphere is an essential part of the carbon cycle. Carbon TOC Study Guide TOC Textbook Website MHHE Website 706 CHEMISTRY IN THE ATMOSPHERE Carbon dioxide in atmosphere Assimilation by plants Plant respiration Animal respiration Litter Dead organisms Decomposition Root respiration FIGURE 17.11 Soil respiration The carbon cycle. dioxide is produced when any form of carbon or a carbon-containing compound is burned in an excess of oxygen. Many carbonates give off CO2 when heated, and all give off CO2 when treated with acid: CaCO3(s) 88n CaO(s) CaCO3(s) CO2(g) 2HCl(aq) 88n CaCl2(aq) H2O(l ) CO2(g) Carbon dioxide is also a by-product of the fermentation of sugar: yeast C6H12O6(aq) 88n 2C2H5OH(aq) glucose ethanol 2CO2(g) Carbohydrates and other complex carbon-containing molecules are consumed by animals, which respire and release CO2 as an end product of metabolism: C6H12O6(aq) 6O2(g) 88n 6CO2(g) 6H2O(l ) As mentioned earlier, another major source of CO2 is volcanic activity. Carbon dioxide is removed from the atmosphere by photosynthetic plants and certain microorganisms: This reaction requires radiant energy (visible light). Back Forward 6CO2(g) 6H2O(l ) 88n C6H12O6(aq) 6O2(g) After plants and animals die, the carbon in their tissues is oxidized to CO2 and returns to the atmosphere. In addition, there is a dynamic equilibrium between atmospheric CO2 and carbonates in the oceans and lakes. Main Menu TOC Study Guide TOC Textbook Website MHHE Website 17.5 FIGURE 17.12 The incoming radiation from the sun and the outgoing radiation from Earth’s surface. THE GREENHOUSE EFFECT 707 Energy Incoming solar radiation Outgoing terrestrial radiation 5000 Stable form Stretched Compressed FIGURE 17.13 Vibrational motion of a diatomic molecule. Chemical bonds can be stretched and compressed like a spring. Back Forward Main Menu 15,000 Wavelength (nm) 25,000 The solar radiant energy received by Earth is distributed over a band of wavelengths between 100 and 5000 nm, but much of it concentrated in the 400- to 700-nm range, which is the visible region of the spectrum (Figure 17.12). By contrast, the thermal radiation emitted by Earth’s surface is characterized by wavelengths longer than 4000 nm (IR region) because of the much lower average surface temperature compared to that of the sun. The outgoing IR radiation can be absorbed by water and carbon dioxide, but not by nitrogen and oxygen. All molecules vibrate, even at the lowest temperatures. The energy associated with molecular vibration is quantized, much like the electronic energies of atoms and molecules. To vibrate more energetically, a molecule must absorb a photon of a specific wavelength in the IR region. First, however, its dipole moment must change during the course of a vibration. [Recall that the dipole moment of a molecule is the product of the charge and the distance between charges (see p. 377).] Figure 17.13 shows how a diatomic molecule can vibrate. If the molecule is homonuclear like N2 and O2, there can be no change in the dipole moment; the molecule has a zero dipole moment no matter how far apart or close together the two atoms are. We call such molecules IRinactive because they cannot absorb IR radiation. On the other hand, all heteronuclear diatomic molecules are IR-active; that is, they all can absorb IR radiation because their dipole moments constantly change as the bond lengths change. A polyatomic molecule can vibrate in more than one way. Water, for example, can vibrate in three different ways, as shown in Figure 17.14. Because water is a polar molecule, it is easy to see that any of these vibrations results in a change in dipole moment because there is a change in bond length. Therefore, a H2O molecule is IRactive. Carbon dioxide has a linear geometry and is nonpolar. Figure 17.15 shows two of the four ways a CO2 molecule can vibrate. One of them [Figure 17.15(a)] symmetrically displaces atoms from the center of gravity and will not create a dipole moment, but the other vibration [Figure 17.15(b)] is IR-active because the dipole moment changes from zero to a maximum value in one direction and then reaches the same maximum value when it changes to the other extreme position. TOC Study Guide TOC Textbook Website MHHE Website 708 CHEMISTRY IN THE ATMOSPHERE (a) (b) FIGURE 17.15 Two of the four ways a carbon dioxide molecule can vibrate. The vibration in (a) does not result in a change in dipole moment, but the vibration in (b) renders the molecule IR-active. Upon receiving a photon in the IR region, a molecule of H2O or CO2 is promoted to a higher vibrational energy level: H2O FIGURE 17.14 The three different modes of vibration of a water molecule. Each mode of vibration can be imagined by moving the atoms along the arrows and then reversing their directions. h 88n H2O* CO2 h 88n CO2* (the asterisk denotes a vibrationally excited molecule). These energetically excited molecules soon lose their excess energy either by collision with other molecules or by spontaneous emission of radiation. Part of this radiation is emitted to outer space and part returns to Earth’s surface. Although the total amount of water vapor in our atmosphere has not altered noticeably over the years, the concentration of CO2 has been rising steadily since the turn of the century as a result of the burning of fossil fuels (petroleum, natural gas, and coal). Figure 17.16 shows the percentages of CO2 emitted due to human activities in the United States in 1995, and Figure 17.17 shows the variation of carbon dioxide concentration over a period of years, as measured in Hawaii. In the Northern Hemisphere, the seasonal oscillations are caused by removal of carbon dioxide by photosynthesis during the growing season and its buildup during the fall and winter months. Clearly, the trend is toward an increase in CO2. The current rate of increase is about 1 ppm (1 part CO2 per million parts air) by volume per year, which is equivalent to 9 109 tons of CO2! Scientists have estimated that by the year 2000 the CO2 concentration will exceed preindustrial levels by about 25 percent. In addition to CO2 and H2O, other greenhouse gases, such as the CFCs, CH4, and NOx, also contribute appreciably to the warming of the atmosphere. Figure 17.18 shows FIGURE 17.16 Sources of carbon dioxide emission in the United States. Note that not all of the emitted CO2 enters the atmosphere. Some of it is taken up by carbon dioxide “sinks,” such as the ocean. Electricity production 35% Cars and trucks 30% Industry 24% 11% Residential heating Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 17.5 THE GREENHOUSE EFFECT 709 CO2 concentration (ppm by volume) 340 330 320 1960 1965 1970 1975 1980 1984 FIGURE 17.17 Yearly variation of carbon dioxide concentration at Mauna Loa, Hawaii. The general trend clearly points to an increase of carbon dioxide in the atmosphere. This trend has continued: in 1994 the CO2 concentration reached 380 ppm. The difference in global temperatures between today and the last ice age is only 4 – 5°C. the gradual increase in temperature over the years and Figure 17.19 shows the relative contributions of the greenhouse gases to global warming. It is predicted by some meteorologists that should the buildup of greenhouse gases continue at its current rate, Earth’s average temperature will increase by about 1° to 3°C in the twenty-first century. Although a temperature increase of a few degrees may FIGURE 17.18 Temperature increase on Earth’s surface from 1880 to 1996. Zero represents the average temperature for the years 1951–1980. (Source: NASA Goddard Institute for Space Studies.) 0.6 Temperature (°C) 0.4 0.2 0 –0.2 –0.4 –0.6 1880 Back Forward Main Menu TOC 1900 1920 1940 Study Guide TOC 1960 1980 2000 Textbook Website MHHE Website 710 CHEMISTRY IN THE ATMOSPHERE FIGURE 17.19 Contribution to global warming by various greenhouse gases. The concentrations of CFCs and methane are much lower than that of carbon dioxide. However, because they can absorb IR radiation much more effectively than CO2, they make an appreciable contribution to the overall warming effect. As more nations industralize, the production of CO2 will increase appreciably. CO2 55% N2O 6% CH4 15% CFCs 24% seem insignificant, it is actually large enough to disrupt the delicate thermal balance on Earth and could cause glaciers and icecaps to melt. Consequently, the sea level would rise and coastal areas would be flooded. Of course, predicting weather trends is extremely difficult, and there are other potentially moderating factors to take into account before concluding that global warming is inevitable and irreversible. For example, the ash from volcanic eruptions diffuses upward and can stay in the atmosphere for years. By reflecting incoming sunlight, volcanic ash can cause a cooling effect. Furthermore, the warming effect of CFCs in the troposphere is offset by its action in the stratosphere. Since ozone is a polar polyatomic molecule, it is also an effective greenhouse gas. A decrease in ozone brought about by CFCs actually produces a noticeable drop in temperature. To combat the greenhouse effect, if it indeed is a serious environmental problem, we must lower carbon dioxide emission. This can be done by improving energy efficiency in automobiles and in household heating and lighting, and by developing nonfossil fuel energy sources, such as photovoltaic cells. Nuclear energy is a viable alternative, but its use is highly controversial due to the difficulty of disposing of radioactive waste and the fact that nuclear power stations are more prone to accidents than conventional power stations (see Chapter 23). The proposed phasing out of CFCs, the most potent greenhouse gas, will help to slow down the warming trend. The recovery of methane gas generated at landfills and the reduction of natural gas leakages are other steps we could take to control CO2 emission. Finally, the preservation of the Amazon jungle, tropical forests in Southeast Asia, and other large forests is vital to maintaining the steady-state concentration of CO2 in the atmosphere. Converting forests to farmland for crops and grassland for cattle may do irreparable damage to the delicate ecosystem and permanently alter the climate pattern on Earth. EXAMPLE 17.2 Which of the following gases qualify as a greenhouse gas: CO, NO, NO2, Cl2, H2, Ne? Similar problem: 17.36. Answer Only CO, NO, and NO2, which are all polar molecules, qualify as greenhouse gases. Both Cl2 and H2 are homonuclear diatomic molecules, and Ne is atomic. These three species are all IR-inactive. PRACTICE EXERCISE Which of the following is a more effective greenhouse gas: CO or H2O? Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 17.6 17.6 Scientists have known about acid rain since the late nineteenth century, but it has been a public issue for only about 25 years. ACID RAIN 711 ACID RAIN Every year acid rain causes hundreds of millions of dollars’ worth of damage to stone buildings and statues throughout the world. The term “stone leprosy” is used by some environmental chemists to describe the corrosion of stone by acid rain (Figure 17.20). Acid rain is also toxic to vegetation and aquatic life. Many well-documented cases show dramatically how acid rain has destroyed agricultural and forest lands and killed aquatic organisms (see Figure 15.10). Precipitation in the northeastern United States has an average pH of about 4.3 (Figure 17.21). Since atmospheric CO2 in equilibrium with rainwater would not be expected to result in a pH less than 5.5, sulfur dioxide (SO2) and, to a lesser extent, nitrogen oxides from auto emissions are believed to be responsible for the high acidity of rainwater. Acidic oxides, such as SO2, react with water to give the corresponding acids. There are several sources of atmospheric SO2. Nature itself contributes much SO2 in the form of volcanic eruptions. Also, many metals exist combined with sulfur in nature. Extracting the metals often entails smelting, or roasting, the ores — that is, heating the metal sulfide in air to form the metal oxide and SO2. For example, 2ZnS(s) 3O2(g) 88n 2ZnO(s) 2SO2(g) The metal oxide can be reduced more easily than the sulfide (by a more reactive metal or in some cases by carbon) to the free metal. Although smelting is a major source of SO2, the burning of fossil fuels in industry, in power plants, and in homes accounts for most of the SO2 emitted to the atmosphere (Figure 17.22). The sulfur content of coal ranges from 0.5 to 5 percent by mass, depending on the source of the coal. The sulfur content of other fossil fuels is similarly variable. Oil from the Middle East, for instance, is low in sulfur, while that from Venezuela has a high sulfur content. To a lesser extent, the nitrogen-containing compounds in oil and coal are converted to nitrogen oxides, which can also acidify rainwater. FIGURE 17.20 Photos of a statue, taken about 60 years apart (in 1908 and 1969), show the damaging effects of air pollutants such as SO2. Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 712 CHEMISTRY IN THE ATMOSPHERE FIGURE 17.21 Mean precipitation pH in the United States in 1994. Most SO2 comes from the midwestern states. Prevailing winds carry the acid droplets formed over the Northeast. Nitrogen oxides also contribute to the acid rain formation. 5.3 5.3 5.1 4.9 5.3 4.5 4.7 4.5 4.3 4.3 4.5 4.5 5.3 5.1 5.1 4.7 4.9 5.1 All in all, some 50 million to 60 million tons of SO2 are released into the atmosphere each year! Some of it is oxidized to SO3. For example, it may react with ozone in the troposphere as follows: SO2(g) O3(g) 88n SO3(g) O2(g) (The ozone that takes part in this reaction is formed by the action of sunlight on nitrogen oxides and hydrocarbons. See the next section.) Or a SO2 molecule may be photoexcited by sunlight: SO2(g) h 88n SO2*(g) and then undergo the reactions FIGURE 17.22 Sulfur dioxide and other air pollutants being released into the atmosphere from a coal-burning power plant. Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 17.6 SO2*(g) 713 O2(g) 88n SO4(g) SO4(g) ACID RAIN SO2(g) 88n 2SO3(g) (The neutral SO4 species is a reactive intermediate.) In another reaction sequence, solid particles in the atmosphere can act as heterogeneous catalysts for the reaction 2SO2(g) SO2 reacts with water to form H and HSO3 . O2(g) 88n 2SO3(g) The SO3 is converted to H2SO4 by rainwater. In this form it can corrode limestone and marble (CaCO3). A typical reaction is CaCO3(s) H2SO4(aq) 88n CaSO4(s) H2O(l ) CO2(g) Sulfur dioxide can also attack calcium carbonate directly: 2CaCO3(s) 2SO2(g) O2(g) 88n 2CaSO4(s) 2CO2(g) There are two ways to minimize the effects of SO2 pollution. The most direct approach is to remove sulfur from fossil fuels before combustion, but this is technologically difficult to accomplish. A cheaper but less efficient way is to remove SO2 as it is formed. For example, in one process powdered limestone is injected into the power plant boiler or furnace along with coal (Figure 17.23). At high temperatures the following decomposition occurs: FIGURE 17.23 Common procedure for removing SO2 from burning fossil fuel. Powdered limestone decomposes into CaO, which reacts with SO2 to form CaSO3. The remaining SO2 is reacted with an aqueous suspension of CaO to form CaSO3. QQQQ ;;;; QQQQ ;;;; QQQQ ;;;; QQQQ ;;;; QQQQ ;;;; QQQQ ;;;; QQQQ ;;;; Mostly CO2 and air Smokestack Furnace CaCO3 S + O2 CaCO3 CaO + SO2 Purification chamber Aqueous suspension of CaO SO2 CaO + CO2 CaSO3 SO2, CO2 Air Air Coal Back Forward Main Menu CaSO3 TOC Study Guide TOC Textbook Website MHHE Website 714 CHEMISTRY IN THE ATMOSPHERE FIGURE 17.24 Spreading calcium oxide (CaO) over acidified soil. This process is called liming. CaCO3(s) 88n CaO(s) CO2(g) limestone quicklime The quicklime reacts with SO2 to form calcium sulfite and some calcium sulfate: CaO(s) 2CaO(s) SO2(g) 88n CaSO3(s) 2SO2(g) O2(g) 88n 2CaSO4(s) To remove any remaining SO2, an aqueous suspension of quicklime is injected into a purification chamber prior to the gases’ escape through the smokestack. Quicklime is also added to lakes and soils in a process called liming to reduce their acidity (Figure 17.24). Installing a sulfuric acid plant near a metal ore refining site is also an effective way to cut SO2 emission because the SO2 produced by roasting metal sulfides can be captured for use in the synthesis of sulfuric acid. This is a very sensible way to turn what is a pollutant in one process into a starting material for another process! 17.7 PHOTOCHEMICAL SMOG The word “smog” was coined to describe the combination of smoke and fog that shrouded London during the 1950s. The primary cause of this noxious cloud was sulfur dioxide. Today, however, we are more familiar with photochemical smog which is formed by the reactions of automobile exhaust in the presence of sunlight. Automobile exhaust consists mainly of NO, CO, and various unburned hydrocarbons. These gases are called primary pollutants because they set in motion a series of photochemical reactions that produce secondary pollutants. It is the secondary pollutants—chiefly NO2 and O3—that are responsible for the buildup of smog. Nitric oxide is the product of the reaction between atmospheric nitrogen and oxygen at high temperatures inside an automobile engine: The heavy use of automobiles is the cause of photochemical smog formation. Back Forward N2(g) O2(g) 88n 2NO(g) Once released into the atmosphere, nitric oxide is quickly oxidized to nitrogen dioxide: Main Menu TOC Study Guide TOC Textbook Website MHHE Website 17.7 2NO(g) PHOTOCHEMICAL SMOG 715 O2(g) 88n 2NO2(g) Sunlight causes the photochemical decomposition of NO2 (at a wavelength shorter than 400 nm) into NO and O: NO2(g) h 88n NO(g) O(g) Atomic oxygen is a highly reactive species that can initiate a number of important reactions, one of which is the formation of ozone: O(g) Ozone plays a dual role in the atmosphere: It is Dr. Jekyll in the stratosphere and Mr. Hyde in the troposphere. M 88n O3(g) M where M is some inert substance such as N2. Ozone attacks the CPC linkage in rubber: R R An aldehyde is an organic compound containing the OCHO group. The simplest member is formaldehyde, HCHO. O2(g) G D CP C D G R R R O G D C C D G ROOR O3 R R H2O R G C PO D D OPC G R H2O2 R where R represents groups of C and H atoms. In smog-ridden areas, this reaction can cause automobile tires to crack. Similar reactions are also damaging to lung tissues and other biological substances. Ozone can be formed also by a series of very complex reactions involving unburned hydrocarbons, aldehydes, nitrogen oxides, and oxygen. One of the products of these reactions is peroxyacetyl nitrate, PAN: H3CO COOOOONO2 B O PAN is a powerful lachrymator, or tear producer, and causes breathing difficulties. Figure 17.25 shows typical variations with time of primary and secondary pollutants. Initially, the concentration of NO2 is quite low. As soon as solar radiation penetrates the atmosphere, more NO2 is formed from NO and O2. Note that the concentration of ozone remains fairly constant at a low level in the early morning hours. As the concentration of unburned hydrocarbons and aldehydes increases in the air, the concentrations of NO2 and O3 also rise rapidly. The actual amounts, of course, depend on the location, traffic, and weather conditions, but their presence is always accompanied by haze (Figure 17.26). The oxidation of hydrocarbons produces various organic intermediates, such as alcohols and carboxylic acids, which are all less volatile than the Relative concentrations FIGURE 17.25 Typical variations with time in concentration of air pollutants on a smoggy day. Hydrocarbons NO2 O3 NO 4 6 8 A.M. Back Forward Main Menu TOC 10 12 Noon Study Guide TOC 2 4 6 P.M. Textbook Website MHHE Website 716 CHEMISTRY IN THE ATMOSPHERE FIGURE 17.26 A smoggy day in New York City. The haze over the Smoky Mountains is caused by aerosols produced by the oxidation of hydrocarbons emitted by pine trees. hydrocarbons themselves. These substances eventually condense into small droplets of liquid. The dispersion of these droplets in air, called aerosol, scatters sunlight and reduces visibility. This interaction also makes the air look hazy. As the mechanism of photochemical smog formation has become better understood, major efforts have been made to reduce the buildup of primary pollutants. Most automobiles now are equipped with catalytic converters designed to oxidize CO and unburned hydrocarbons to CO2 and H2O and to reduce NO and NO2 to N2 and O2 (see Section 13.6). More efficient automobile engines and better public transportation systems would also help to decrease air pollution in urban areas. A recent technological innovation to combat photochemical smog is to coat automobile radiators and air conditioner compressors with a platinum catalyst. So equipped, a running car can purify the air that flows under the hood by converting ozone and carbon monoxide to oxygen and carbon dioxide: O3(g) Pt CO(g) 88n O2(g) CO2(g) In a city like Los Angeles, where the number of miles driven in one day equals nearly 300 million, this approach would significantly improve the air quality and reduce the “high-ozone level” warnings frequently issued to its residents. In fact, a drive on the freeway would help to clean up the air! 17.8 INDOOR POLLUTION Difficult as it is to avoid air pollution outdoors, it is no easier to avoid indoor pollution. The air quality in homes and in the workplace is affected by human activities, by construction materials, and by other factors in our immediate environment. The common indoor pollutants are radon, carbon monoxide and carbon dioxide, and formaldehyde. THE RISK FROM RADON In a highly publicized case about 25 years ago, an employee reporting for work at a nuclear power plant in Pennsylvania set off the plant’s radiation monitor. Astonishingly, the source of his contamination turned out not to be the plant, but radon in his home! A lot has been said and written about the potential dangers of radon as an air pollutant. Just what is radon? Where does it come from? And how does it affect our health? Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 17.8 INDOOR POLLUTION 717 FIGURE 17.27 Sources of background radiation. Radon 54% Other man–made radiation 3% Nuclear testing 2% The uranium decay series is discussed in Chapter 23. Activated charcoal is heat-treated charcoal whose surface has a greater tendency to bind molecules. Medical procedures 14% Other natural radiation 27% Radon is a member of Group 8A (the noble gases). It is an intermediate product of the radioactive decay of uranium-238. All isotopes of radon are radioactive, but radon-222 is the most hazardous because it has the longest half-life—3.8 days. Radon, which accounts for slightly over half the background radioactivity on Earth, is generated mostly from the phosphate minerals of uranium (Figure 17.27). Since the 1970s, high levels of radon have been detected in homes built on reclaimed land above uranium mill tailing deposits. The colorless, odorless, and tasteless radon gas enters a building through tiny cracks in the basement floor. It is slightly soluble in water, so it can be spread in different media. Radon-222 is an -emitter. When it decays, it produces radioactive polonium-214 and polonium-218, which can build up to high levels in an enclosed space. These solid radioactive particles can adhere to airborne dust and smoke, which are inhaled into the lungs and deposited in the respiratory tract. Over a long period of time, the particles emitted by polonium and its decay products, which are also radioactive, can cause lung cancer. What can be done to combat radon pollution indoors? The first step is to measure the radon level in the basement with a reliable test kit. Short-term and long-term kits are available (Figure 17.28). The short-term tests use activated charcoal to collect the decay products of radon over a period of several days. The container is sent to a laboratory where a technician measures the radioactivity ( rays) from radon-decay products lead-214 and bismuth-214. Knowing the length of exposure, the lab technician back-calculates to determine radon concentration. The long-term test kits use a piece of special polymer film on which an particle will leave a “track.” After several months’ exposure, the film is etched with a sodium hydroxide solution and the num- FIGURE 17.28 Home radon detectors: Long-term track etch (left) and short-term charcoal canister (right). Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website 718 CHEMISTRY IN THE ATMOSPHERE ber of tracks counted. Knowing the length of exposure enables the technician to calculate the radon concentration. If the radon level is unacceptably high, then the house must be regularly ventilated. This precaution is particularly important in recently built houses, which are well insulated. A more effective way to prevent radon pollution is to reroute the gas before it gets into the house, for example, by installing a ventilation duct to draw air from beneath the basement floor to the outside. Currently there is considerable controversy regarding the health effects of radon. The first detailed studies of the effects of radon on human health were carried out in the 1950s when it was recognized that uranium miners suffered from an abnormally high incidence of lung cancer. Some scientists have challenged the validity of these studies because the miners were also smokers. It seems quite likely that there is a synergistic effect between radon and smoking on the development of lung cancer. Radondecay products will adhere not only to tobacco tar deposits in the lungs, but also to the solid particles in cigarette smoke, which can be inhaled by smokers and nonsmokers. More systematic studies are needed to evaluate the environmental impact of radon. In the meantime, the Environmental Protection Agency (EPA) has recommended remedial action where the radioactivity level due to radon exceeds 4 picocuries (pCi) per liter of air. (A curie corresponds to 3.70 1010 disintegrations of radioactive nuclei per second; a picocurie is a trillionth of a curie, or 3.70 10 2 disintegrations per second.) EXAMPLE 17.3 Half-life is defined in Section 13.3. Recall that radioactive decays obey first-order kinetics. The half-life of Rn-222 is 3.8 days. Starting with 1.0 g of Rn-222, how much will be left after 10 half-lives? We can generalize the fraction of the isotope left after n half-lives by ( 1 )n. 2 Setting n 10, we write Answer quantity of Rn-222 left 1 10 2 1.0 g 9.8 10 4 g PRACTICE EXERCISE The concentration of Rn-222 in the basement of a house is 1.8 10 6 mol/L. Assume the air remains static and calculate the concentration of the radon after 2.4 days. CARBON DIOXIDE AND CARBON MONOXIDE Both carbon dioxide (CO2) and carbon monoxide (CO) are products of combustion. In the presence of an abundant supply of oxygen, CO2 is formed; in a limited supply of oxygen, both CO and CO2 are formed. The indoor sources of these gases are gas cooking ranges, woodstoves, space heaters, tobacco smoke, human respiration, and exhaust fumes from cars (in garages). Carbon dioxide is not a toxic gas, but it does have an asphyxiating effect (see Chemistry in Action on p. 480). In airtight buildings, the concentration of CO2 can reach as high as 2000 ppm by volume (compared with 3 ppm outdoors). Workers exposed to high concentrations of CO2 in skyscrapers and other sealed environments become fatigued more easily and have difficulty concentrating. Adequate ventilation is the solution to CO2 pollution. Back Forward Main Menu TOC Study Guide TOC Textbook Website MHHE Website SUMMARY OF FACTS AND CONCEPTS The concentration of carboxyhemoglobin in the blood of regular smokers is two to five times higher than that in nonsmokers. 719 Like CO2, CO is a colorless and odorless gas, but it differs from CO2 in that it is highly poisonous. The toxicity of CO lies in its unusual ability to bind very strongly to hemoglobin, the oxygen carrier in blood. Both O2 and CO bind to the Fe(II) ion in hemoglobin, but the affinity of hemoglobin for CO is about 200 times greater than that for O2 (see Chapter 25). Hemoglobin molecules with tightly bound CO (called carboxyhemoglobin) cannot carry the oxygen needed for metabolic processes. A small amount of CO intake can cause drowsiness and headache; death may result when about half the hemoglobin molecules are complexed with CO. The best first-aid response to CO poisoning is to remove the victim immediately to an area with a plentiful oxygen supply or to give mouth-to-mouth resuscitation. FORMALDEHYDE Formaldehyde (HCHO) is a rather disagreeable-smelling liquid used as a preservative for laboratory specimens. Industrially, formaldehyde resins are used as bonding agents in building and furniture construction materials such as plywood and particle board. In addition, urea-formaldehyde insulation foams are used to fill wall cavities. The resins and foams slowly break down to release free formaldehyde, especially under acid and humid conditions. Low concentrations of formaldehyde in the air can cause drowsiness, nausea, headaches, and other respiratory ailments. Laboratory tests show that breathing high concentrations of formaldehyde can induce cancers in animals, but whether it has a similar effect in humans is unclear. The safe standard of formaldehyde in indoor air has been set at 0.1 ppm by volume. Since formaldehyde is a reducing agent, devices have been constructed to remove it by means of a redox reaction. Indoor air is circulated through an air purifier containing an oxidant such as Al2O3/KMnO4, which converts formaldehyde to the less harmful and less volatile formic acid (HCOOH). Proper ventilation is the best way to remove formaldehyde. However, care should be taken not to remove the air from a room too quickly without replenishment, since a reduced pressure would cause the formaldehyde resins to decompose faster, resulting in the release of more formaldehyde. SUMMARY OF FACTS AND CONCEPTS Back Forward Main Menu 1. Earth’s atmosphere is made up mainly of nitrogen and oxygen, plus a number of other trace gases. The chemical processes that go on in the atmosphere are influenced by solar radiation, volcanic eruption, and human activities. 2. In the outer regions of the atmosphere the bombardment of molecules and atoms by solar particles gives rise to auroras. The glow on space shuttles is caused by excitation of molecules adsorbed on the shuttles’ surface. 3. Ozone in the stratosphere absorbs harmful UV radiation in the 200- to 300-nm range and protects life underneath. For many years, chlorofluorocarbons have been destroying the ozone layer. 4. Volcanic eruptions can lead to air pollution, deplete ozone in the stratosphere, and affect climate. 5. Carbon dioxide’s ability to absorb infrared radiation enables it to trap some of the outgoing heat from Earth, warming its surface. Other gases such as the CFCs and methane also contribute to global warming. 6. Sulfur dioxide, and to a lesser extent nitrogen oxides, generated mainly from the burning of fossil fuels and from the roasting of metal sulfides, causes acid rain. 7. Photochemical smog is formed by the photochemical reaction of automobile exhaust in the TOC Study Guide TOC Textbook Website MHHE Website 720 CHEMISTRY IN THE ATMOSPHERE presence of sunlight. It is a complex reaction involving nitrogen oxides, ozone, and hydrocarbons. 8. Indoor air pollution is caused by radon, a radioactive gas formed during uranium decay; carbon monoxide and carbon dioxide, products of combustion; and formaldehyde, a volatile organic substance released from resins used in construction materials. KEY WORDS Greenhouse effect, p. 705 Ionosphere, p. 695 Mesosphere, p. 695 Nitrogen fixation, p. 694 Photochemical smog, p. 614 Stratosphere, p. 695 Thermosphere, p. 695 Troposphere, p. 694 QUESTIONS AND PROBLEMS EARTH’S ATMOSPHERE Review Questions 17.1 Describe the regions of Earth’s atmosphere. 17.2 Briefly outline the main processes of the nitrogen and oxygen cycles. 17.3 Explain why, for maximum performance, supersonic airplanes need to fly at a high altitude (in the stratosphere). 17.4 Jupiter ’s atmosphere consists mainly of hydrogen (90 percent) and helium (9 percent). How does this mixture of gases contrast with the composition of Earth’s atmosphere? Why does the composition differ? ergy for the oxygen-to-hydrogen bond in OH as 460 kJ. What is the longest wavelength (in nm) of radiation that can bring about the reaction OH(g) 88n O(g) H(g) 17.12 The green color observed in aurora borealis is produced by the emission of a photon by an electronically excited oxygen atom at 558 nm. Calculate the energy difference between the two levels involved in the emission process. DEPLETION OF OZONE IN THE STRATOSPHERE Review Questions Problems 17.5 Referring to Table 17.1, calculate the mole fraction of CO2 and its concentration in parts per million by volume. 17.6 Calculate the partial pressure of CO2 (in atm) in dry air when the atmospheric pressure is 754 mmHg. 17.7 Describe the processes that result in the warming of the stratosphere. 17.8 Calculate the total mass (in kilograms) of nitrogen, oxygen, and carbon dioxide gases in the atmosphere. (Hint: See Problem 5.92 and Table 17.1. Use 29.0 g/mol for the molar mass of air.) PHENOMENA IN THE OUTER LAYERS OF THE ATMOSPHERE 17.13 Briefly describe the absorption of solar radiation in the stratosphere by O2 and O3 molecules. 17.14 Explain the processes that have a warming effect on the stratosphere. 17.15 List the properties of CFCs, and name four major uses of these compounds. 17.16 How do CFCs and nitrogen oxides destroy ozone in the stratosphere? 17.17 What causes the polar ozone holes? 17.18 How do volcanic eruptions contribute to ozone destruction? 17.19 Describe ways to curb the destruction of ozone in the stratosphere. 17.20 Discuss the effectiveness of some of the CFC substitutes. Review Questions 17.9 What process gives rise to aurora borealis and aurora australis? 17.10 Why can astronauts not release oxygen atoms to test the mechanism of shuttle glow? Problems 17.11 The highly reactive OH radical (a species with an unpaired electron) is believed to be involved in some atmospheric processes. Table 9.4 lists the bond en- Back Forward Main Menu TOC Problems 17.21 Given that the quantity of ozone in the stratosphere is equivalent to a 3.0-mm-thick layer of ozone on Earth at STP, calculate the number of ozone molecules in the stratosphere and their mass in kilograms. (Hint: The radius of Earth is 6371 km and the surface area of a sphere is 4 r 2, where r is the radius.) 17.22 Referring to the answer in Problem 17.21, and assuming that the level of ozone in the stratosphere has Study Guide TOC Textbook Website MHHE Website QUESTIONS AND PROBLEMS 17.23 17.24 17.25 17.26 already fallen 6.0 percent, calculate the number of kilograms of ozone that would have to be manufactured on a daily basis so that we could restore the ozone to the original level in 100 years. If ozone is made according to the process 3O2(g) 88n 2O3(g), how many kilojoules of energy would be required? Both Freon-11 and Freon-12 are made by the reaction of carbon tetrachloride (CCl4) with hydrogen fluoride. Write equations for these reactions. Why are CFCs not decomposed by UV radiation in the troposphere? The average bond energies of the COCl and COF bonds are 340 kJ/mol and 485 kJ/mol, respectively. Based on this information, explain why the COCl bond in a CFC molecule is preferentially broken by solar radiation at 250 nm. Like CFCs, certain bromine-containing compounds such as CF3Br can also participate in the destruction of ozone by a similar mechanism starting with the Br atom: CF3Br 88n CF3 Br Given that the average COBr bond energy is 276 kJ/mol, estimate the longest wavelength required to break this bond. Will this compound be decomposed in the troposphere only or in both the troposphere and stratosphere? 17.27 Draw Lewis structures for chlorine nitrate (ClONO2) and chlorine monoxide (ClO). 17.28 Draw Lewis structures for HCFC-123 (CF3CHCl2) and CF3CFH2. VOLCANOES Review Questions 17.29 What are the effects of volcanic eruptions on climate? 17.30 Classify the reaction between H2S and SO2 that leads to the formation of sulfur at the site of a volcanic eruption. THE GREENHOUSE EFFECT Review Questions 17.31 What is the greenhouse effect? What is the criterion for classifying a gas as a greenhouse gas? 17.32 Why is more emphasis placed on the role of carbon dioxide in the greenhouse effect than on that of water? 17.33 Describe three human activities that generate carbon dioxide. List two major mechanisms for the uptake of carbon dioxide. 17.34 Deforestation contributes to the greenhouse effect in two ways. What are they? 17.35 How does an increase in world population enhance the greenhouse effect? Back Forward Main Menu TOC 721 17.36 Is ozone a greenhouse gas? If so, sketch three ways an ozone molecule can vibrate. 17.37 What effects do CFCs and their substitutes have on Earth’s temperature? 17.38 Why are CFCs more effective greenhouse gases than methane and carbon dioxide? Problems 17.39 The annual production of zinc sulfide (ZnS) is 4.0 104 tons. Estimate the number of tons of SO2 produced by roasting it to extract zinc metal. 17.40 Calcium oxide or quicklime (CaO) is used in steelmaking, cement manufacture, and pollution control. It is prepared by the thermal decomposition of calcium carbonate: CaCO3(s) 88n CaO(s) CO2(g) Calculate the yearly release of CO2 (in kilograms) to the atmosphere if the annual production of CaO in the United States is 1.7 1010 kg. 17.41 The molar heat capacity of a diatomic molecule is 29.1 J/K mol. Assuming the atmosphere contains only nitrogen gas and there is no heat loss, calculate the total heat intake (in kilojoules) if the atmosphere warms up by 3°C during the next 50 years. Given that there are 1.8 1020 moles of diatomic molecules present, how many kilograms of ice (at the North and South Poles) will this quantity of heat melt at 0°C? (The molar heat of fusion of ice is 6.01 kJ/mol.) 17.42 As mentioned in the chapter, spraying the stratosphere with hydrocarbons such as ethane and propane should eliminate Cl atoms. What is the drawback of this procedure if used on a large scale for an extended period of time? ACID RAIN Review Questions 17.43 Name the gas that is largely responsible for the acid rain phenomenon. 17.44 List three detrimental effects of acid rain. 17.45 Briefly discuss two industrial processes that lead to acid rain. 17.46 Discuss ways to curb acid rain. 17.47 Water and sulfur dioxide are both polar molecules and their geometry is similar. Why is SO2 not considered a major greenhouse gas? 17.48 Describe the removal of SO2 by CaO (to form CaSO3) in terms of a Lewis acid-base reaction. Problems 17.49 An electric power station annually burns 3.1 107 kg of coal containing 2.4 percent sulfur by mass. Calculate the volume of SO2 emitted at STP. Study Guide TOC Textbook Website MHHE Website 722 CHEMISTRY IN THE ATMOSPHERE 17.50 The concentration of SO2 in the troposphere over a certain region is 0.16 ppm by volume. The gas dissolves in rainwater as follows: SO2(g) H2O(l ) 34 H (aq) at noon in June, (b) New York City at 1 p.m. in July, (c) Boston at noon in January. Explain your choice. INDOOR POLLUTION HSO3 (aq) Review Questions Given that the equilibrium constant for the above reaction is 1.3 10 2, calculate the pH of the rainwater. Assume that the reaction does not affect the partial pressure of SO2. PHOTOCHEMICAL SMOG 17.61 List the major indoor pollutants and their sources. 17.62 What is the best way to deal with indoor pollution? 17.63 Why is it dangerous to idle a car ’s engine in a poorly ventilated place, such as the garage? 17.64 Describe the properties that make radon an indoor pollutant. Would radon be more hazardous if 222Rn had a longer half-life? Review Questions 17.51 What is photochemical smog? List the factors that favor the formation of photochemical smog. 17.52 What are primary and secondary pollutants? 17.53 Identify the gas that is responsible for the brown color of photochemical smog. 17.54 The safety limits of ozone and carbon monoxide are 120 ppb by volume and 9 ppm by volume, respectively. Why does ozone have a lower limit? 17.55 Suggest ways to minimize the formation of photochemical smog. 17.56 In which region of the atmosphere is ozone beneficial? In which region is it detrimental? Problems O2(g) 88n 2NO2(g) is an elementary reaction. (a) Write the rate law for this reaction. (b) A sample of air at a certain temperature is contaminated with 2.0 ppm of NO by volume. Under these conditions, can the rate law be simplified? If so, write the simplified rate law. (c) Under the conditions described in (b), the half-life of the reaction has been estimated to be 6.4 103 min. What would the half-life be if the initial concentration of NO were 10 ppm? 17.58 The gas-phase decomposition of peroxyacetyl nitrate (PAN) obeys first-order kinetics: CH3COOONO2 88n CH3COOO 4 NO2 1 Forward Main Menu TOC 17.67 Briefly describe the harmful effects of the following substances: O3, SO2, NO2, CO, CH3COOONO2 (PAN), Rn. 17.68 The equilibrium constant (KP) for the reaction N2(g) O2(g) 34 2NO(g) 31 is 4.0 10 at 25°C and 2.6 10 6 at 1100°C, the temperature of a running car ’s engine. Is this an endothermic or exothermic reaction? 17.69 As stated in the chapter, carbon monoxide has a much higher affinity for hemoglobin than oxygen does. (a) Write the equilibrium constant expression (Kc) for the following process: CO(g) with a rate constant of 4.9 10 s . Calculate the rate of decomposition in M/s if the concentration of PAN is 0.55 ppm by volume. Assume STP conditions. 17.59 On a smoggy day in a certain city the ozone concentration was 0.42 ppm by volume. Calculate the partial pressure of ozone (in atm) and the number of ozone molecules per liter of air if the temperature and pressure were 20.0°C and 748 mmHg, respectively. 17.60 Which of the following settings is the most suitable for photochemical smog formation? (a) Gobi desert Back 17.65 A concentration of 8.00 102 ppm by volume of CO is considered lethal to humans. Calculate the minimum mass of CO in grams that would become a lethal concentration in a closed room 17.6 m long, 8.80 m wide, and 2.64 m high. The temperature and pressure are 20.0°C and 756 mmHg, respectively. 17.66 A volume of 5.0 liters of polluted air at 18.0°C and 747 mmHg is passed through lime water [an aqueous suspension of Ca(OH)2], so that all the carbon dioxide present is precipitated as CaCO3. If the mass of the CaCO3 precipitate is 0.026 g, calculate the percentage by volume of CO2 in the air sample. ADDITIONAL PROBLEMS 17.57 Assume that the formation of nitrogen dioxide: 2NO(g) Problems HbO2(aq) 34 O2(g) HbCO(aq) where HbO2 and HbCO are oxygenated hemoglobin and carboxyhemoglobin, respectively. (b) The composition of a breath of air inhaled by a person smoking a cigarette is 1.9 10 6 mol/L CO and 8.6 10 3 mol/L O2. Calculate the ratio of [HbCO] to [HbO2], given that Kc is 212 at 37°C. 17.70 Instead of monitoring carbon dioxide, suggest another gas that scientists could study to substantiate the fact that CO2 concentration is steadily increasing in the atmosphere. Study Guide TOC Textbook Website MHHE Website QUESTIONS AND PROBLEMS 17.71 In 1991 it was discovered that nitrous oxide (N2O) is produced in the synthesis of nylon. This compound, which is released into the atmosphere, contributes both to the depletion of ozone in the stratosphere and to the greenhouse effect. (a) Write equations representing the reactions between N2O and oxygen atoms in the stratosphere to produce nitric oxide, which is then oxidized by ozone to form nitrogen dioxide. (b) Is N2O a more effective greenhouse gas than carbon dioxide? Explain. (c) One of the intermediates in nylon manufacture is adipic acid [HOOC(CH2)4COOH]. About 2.2 109 kg of adipic acid are consumed every year. It is estimated that for every mole of adipic acid produced, 1 mole of N2O is generated. What is the maximum number of moles of O3 that can be destroyed as a result of this process per year? 17.72 A glass of water initially at pH 7.0 is exposed to dry air at sea level at 20°C. Calculate the pH of the water when equilibrium is reached between atmospheric CO2 and CO2 dissolved in the water, given that Henry’s law constant for CO2 at 20°C is 0.032 mol/L atm. (Hint: Assume no loss of water due to evaporation and use Table 17.1 to calculate the partial pressure of CO2. Your answer should correspond roughly to the pH of rainwater.) 17.73 A 14-m by 10-m by 3.0-m basement had a high radon content. On the day the basement was sealed off from its surroundings so that no exchange of air could take place, the partial pressure of 222Rn was 1.2 10 6 mmHg. Calculate the number of 222Rn isotopes (t 3.8 d) at the beginning and end of 31 days. Assume STP conditions. 17.74 Ozone in the troposphere is formed by the following steps: 723 where O* denotes an electronically excited atom. (a) Explain why the concentration of OH is so small even though the concentrations of O3 and H2O are quite large in the troposphere. (b) What property makes OH a strong oxidizing agent? (c) The reaction between OH and NO2 contributes to acid rain. Write an equation for this process. (d) The hydroxyl radical can oxidize SO2 to H2SO4. The first step is the formation of a neutral HSO3 species, followed by its reaction with O2 and H2O to form H2SO4 and the hydroperoxyl radical (HO2). write equations for these processes. 17.76 The equilibrium constant (KP) for the reaction 2CO(g) O2(g) 34 2CO2(g) is 1.4 1090 at 25°C. Given this enormous value, why doesn’t CO convert totally to CO2 in the troposphere? 17.77 A person was found dead of carbon monoxide poisoning in a well-insulated cabin. Investigation showed that he had used a blackened bucket to heat water on a butane burner. The burner was found to function properly with no leakage. Explain, with an appropriate equation, the cause of his death. 17.78 The carbon dioxide level in the atmosphere today is often compared with that in preindustrial days. Explain how scientists use tree rings and air trapped in polar ice to arrive at the comparison. 17.79 What is funny about the following cartoon? 1 2 NO2 88n NO O O (1) O2 88n O3 (2) The first step is initiated by the absorption of visible light (NO2 is a brown gas.) Calculate the longest wavelength required for step (1) at 25°C. [Hint: You need to first calculate H and hence E for (1). Next, determine the wavelength for decomposing NO2 from E.] 17.75 Although the hydroxyl radical (OH) is present only in a trace amount in the troposphere, it plays a central role in its chemistry because it is a strong oxidizing agent and can react with many pollutants as well as some CFC substitutes (see p. 704). (a) The hydroxyl radical is formed by the following reactions: 320 nm O3 8888888n O* O* Back Forward 17.80 Calculate the standard enthalpy of formation ( H°) f of ClO from the following bond dissociation energies: Cl2: 242.7 kJ/mol; O2: 498.7 kJ/mol; ClO: 206 kJ/mol. Answers to Practice Exercises: 17.2 H2O. 17.3 1.2 10 6 17.1 1.12 103 nm. mol/L. O2 H2O 8888888n 2OH Main Menu TOC Study Guide TOC Textbook Website MHHE Website ...
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This document was uploaded on 07/27/2009.

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