838 CHAPTER 21 En v ironmental Chemistry I Figure 21 . 10 A volcanic eruption. • Although it is unclear whether the CFCs already releas ed to the atmosphere will eventually result in catastrophic damage to life on Earth, it is conceivable that the depletion of ozone can be slow ed by reducing the availability of Cl atoms. Indeed, some chemists have suggested sending a fleet of planes to s pra y 50,000 tons of ethane (C 2 H 6 ) or propane (C 3 H 8 ) 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 + C 2 H 6 +. HCl + C2HS Cl + C3H8 • 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 strato- sphere from airplanes. Technically this solution is feasible, but it would be enormously costly and it would require the collaboration of many nations. Ha ving discussed the chemistry in the outer regions of Earth's atmosphere, we will focus in Sections 21.4 through 2 1. 8 on events closer to us that is, in the troposphere. Volcanoes Volcanic eruptions, Earth 's most spectacular natural displays of energy, are instrumental in form- ing large parts of Earth's crus t. The upper mantle, immediately under the crust, is nearly molten. A slight increase in heat, s uch 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 21.10). An active volcano emits ga ses, liquid s, and solids. The gases spewed into the atmosphere include primarily Nb CO 2 , HC1 , HF , H 2 S, and water vapor. It is estimated that volcanoes are the source of about two-thirds of the sulfur in the air. On the slopes of Mount St. Helens, which last erupted in 1980, deposits of elemental sulfur are visible near the eruption site. At high tempera- tures, the hydrogen sulfide gas given off by a volcano is oxidized by air: • Some of the SO? is reduced by more H 2 S from the volcano to elemental sulfur and water: The re st of the S0 2 is relea sed into the atmosphere, where it reacts with water to form acid rain (see Section 21.6). The tremendous force of a volcanic eruption carries a sizable amount of gas into the stratosphere. There S0 2 is oxidized to S0 3, which is eventually converted to sulfuric acid aero- sols in a series of complex reactions. In addition to destroying ozone in the stratosphere (see page 837), these aerosols can also affect climate. Because the stratosphere is above the atmo- spheric weather patterns, the aerosol cloud s often persist for more than a year. They absorb solar radiation and thereby cause a drop in temperature at Earth's surface. However, this cool- ing effect is local rather than global, because it depends on the site and frequency of volcanic eruptions. 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 (s ee Table 21.1), 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 atmo- sphere, 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 21.11 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 dioxide is produced SECTION 21.5 The Greenhouse Effect 839 Assimilation by plants Fig u re 21.11 The carbon cycle. Carbon dioxide in atmosphere Root respiration - r I .,. -:.' I . • ~ . • • Plant respiration • Dead . orgamsms Animal respiration Decomposition when any form of carbon or a carbon-containing compound is burned in an excess of oxygen. Many carbonates give off CO 2 when heated, and all give off CO 2 when treated with acid: CaC0 3 (s) CaO(s) + COz(g) CaC0 3 (s) + 2HCI(aq) • CaClz(aq) + HzO(l) + COz(g) Carbon dioxide is also a by-product of the fermentation of sugar: glucose ethanol Carbohydrates and other complex carbon-containing molecules are consumed by animals, which respire and release CO 2 as an end product of metabolis m: • Soil re spiration 840 CHAPTER 21 Environmental Chemistry Figure 21.12 The incoming radiation from the sun and the outgoing radiation from Earth's surface. • Stable form Stretc hed Compressed Figure 21.13 Vibration motion of a diatomic molecule. Chemical bonds can be stretched and compressed like a • spn ng. Figure 21.14 The three different modes of vibration of a water molecule. Each mode of vibration can be imagined by moving the atoms along the alTOW S and then reversing the direction of motion. > on ~ OJ c:: W Incoming sol ar radiation 5000 Outaoina o 0 terrestrial radiation 15,000 Wavelength (nm) 25,000 As mentioned earlier, another major source of CO? is volcanic activity. Carbon dioxide is removed from the atmosphere by photosynthetic plants and certain microorganisms: After plants and animals die, the carbon in their tissues is oxidized to CO 2 and returns to the atmo- sphere. In addition, there is a dynamic equilibrium between atmospheric CO 2 and carbonates in the oceans and lakes. The solar radiant energy received by Earth is distributed over a band of wavelengths between 100 and 5000 nm , but much of it is concentrated in the 400- to 700-nm range, which is the vis- ible region of the spectrum (Figure 2l.12). By contrast, the thermal radiation emitted by Earth's surface is characterized by wavelengths longer than 4000 nm (t he 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 molecu- lar vibration is quantized, much like the electronic energies of atoms and molecules. To vibrate more energeticall y, 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 page 289).] Figure 2l.13 shows how a diatomic molecule can vibrate. Ifthe molecule is homonu- clear like N2 and O?, there can be no change in the dipole moment; the molecule ha s a zero dipole moment no matter how far apart or close together the two atoms are. We call such molecules IR- inactive because they cannot absorb IR radiation. On the other hand, all heteronucIear 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 21.14. Because water is a polar molecule, any of these vibrations results in a change in dipole moment because there is a change in bond length. There- fore, an H?O molecule is IR-active. Carbon dioxide has a linear geometry and is nonpolar. Figure 2l.15 shows two of the four ways a CO 2 molecule can vibrate. One of them [Figure 2l.15(a)] symmetrically displaces atoms from the center of gravity and will not create a dipole moment, but the other vibration [Figure 2l.l5(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. Upon receiving a photon in the IR region, a molecule of H 2 0 or CO 2 is promoted to a higher vibrational energy level: H 2 0 + hv -_. H 2 0 * CO? + hv • COi SECTION 21.5 The Greenhouse Effect 841 (a) (b) (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 CO 2 has been rising steadily since the tum of the century as a result of the burning of fossil fuels (petroleum, natural gas, and coal). Figure 21.16 shows the percent- ages of CO 2 emitted due to human activities in the United States in 1998, and Figure 21.17 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 the removal of carbon dioxide by photosynthesis during the growing season and its buildup during the fall and winter months. The trend is toward an increase in CO 2 , The current rate of increase is about 1 ppm (1 part CO 2 per million parts air) by volume per year, which is equivalent to 9 X 10 9 tons of CO 2 ! Scientists have estimated that by the year 2010 the CO 2 concentration will exceed preindustrial levels by about 25 percent. In addition to CO 2 and H 2 0, other greenhouse gases, such as the CFCs, CH 4 , NO x and N 2 0 , also contribute appreciably to the warming of the atmosphere. Figure 21.18 shows the gradual increase in temperature over the years, and Figure 21.19 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 I ° to 3°C in the twenty-first century, Although a temperature increase of a few degrees may seem insignificant, it is actually large enough to disrupt the delicate thermal balance on Earth and could cause glaciers and ice caps to melt. Consequently, the sea level would rise and coastal areas would be flooded. Predicting weather trends is extremely difficult, though, and there are other potentially moderating factors to 340 ~ " S .E 0 ;> ;>., ,n S 0- 0- 330 ~ " 0 .~ - '" b " " () c: 0 () N 0 320 u 1960 1965 1970 1975 Year Fig u re 21.15 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. Electricity production 35 % Industry 24% Cars and trucks 30% Residential heating ll % Figure 21 .16 Sources of carbon dioxide emission in the United States. Note that not all the emitted CO 2 enters the atmosphere. Some of it is taken up by carbon dioxide "sinks," such as the ocean. , 1980 1984 Figure 21.17 Yearly variation of carbon dioxide concentration at Mauna Loa, Hawaii, the source of the longest running record of atmospheric CO 2 , The general trend clearly points to an increase of carbon dioxide in the atmosphere. This trend has continued: in 2007, the CO 2 concentration reached 380 ppm. 842 CHAPTER 21 Environmental Chemistry Figure 21 .18 Temperature ri se on Earth's surface from 18 80 to the present. Zero repre se nt s the average temperature for the years 1951- 1980. • CH 4 CFCs 15 % 24% Figure 21 .19 Contribution to global warming by various greenhouse gases. The concentrations of CFCs and methane are much lower than that of carbon dioxide. However, becau se they can absorb IR radiation much more effectively than CO 2 , the y make a significant contribution to the overall warming effect. Global climate change is the subject of the Academy Award-winning 2006 documentary, An Inconvenient Truth, presented by former vice president AI Gore. Think About It CO 2 , the best- known greenhouse gas, is nonpolar. It is only necessary for at least one of a molecule's vibrational modes to induce a temporary dipole for it to act as a greenhouse gas. 0.6 0.4 ~ 0.2 u <!) B '" 0 k " E ~ -0.2 -0 .4 - -0 .6 1 r~~~_r~~~ 1880 1900 1920 1940 Year 1960 1980 2000 take into account before concluding that global warming is inevitable and irreversible. For exam- ple, 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. Because ozone is a polar poly atomic 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, we must lower carbon dioxide emissions. 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 sta- tions (see Chapter 20). 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 CO 2 emissions. 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 CO 2 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. Which of the following gases qualify as a greenhouse gas: CO, NO, N0 2 , C1 2 , H 2 , Ne? Strategy To behave as a greenhouse gas, either the molecule must possess a dipole moment or some of its vibrational motion s must generate a te mporary dipole moment. Setup The necessary conditions immediately rule out homonuclear diatomic molecules and atomic • speCIes. Solution Only CO, NO, and N0 2 , which are all polar molecules, qualify as greenhouse gases. Both Cl 2 and H2 are homonuclear diatomic molecules , and Ne is atomic. These three species are all IR-inactive. Practice Problem Which of the following is a more effective greenhouse gas: CO or H 2 0? Checkpoint 21.5 The Greenhouse Effect 21.5.1 Which of the following can act as a greenhouse gas? (Select all that apply.) a) CH 4 b) N2 c) Rn d) 0 3 e) Xe • Acid Rain 21.5.2 The greenhouse effect is a) caused by depletion of stratospheric ozone. b) entirely the result of human activity. c) a natural phenomenon that has been enhanced by human activity. d) the absorption of the sun's energy by molecules and atoms in the upper atmosphere. e) responsible for the aurora borealis and the aurora australis . 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 21.20). 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. Precipitation in the northeastern United States has an average pH of about 4.3 (Figure 21.21). Because atmospheric CO 2 in equilibrium with rainwater would not be expected to result in a pH less than 5.5, sulfur dioxide (S02) 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 SOb react with water to give the corresponding acids. There are several sources of atmospheric S02' Nature itself contributes much S02 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 S02' For example, 2ZnS(s) + 30 2 (g) +. 2ZnO(s) + 2S0 2 (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 SOb the burning of fossil fuels in industry, in power plants, and in homes accounts for most of the S02 emitted to the atmosphere (Figure 21.22). The 5.3 5.3 5.1 4.9 4.7 4.5 4.3 4.3 5.3 5.1 5.1 4.7 4.9 5.1 Figure 21.21 Mean precipitation pH in the United States in 1994. Most S0 2 comes from the midwestern states. Prevailing winds carry the acid droplets formed over the Northeast. Nitrogen oxides also contribute to acid rain formation. SECTION 21.6 Acid Rain Figure 21.20 The effect of acid rain on a marble statue. The photos were taken just a few decades apart. ,;' Multimedia Organic and Biochemistry-oil ref i ning processes. 844 CHAPTER 21 Environmental Chemistry Figure 21.22 Sulfur dioxide an ~ other air pollutants being released into the atmosphere from a coal-burning power plant. Figure 21.23 Common procedure for removing S0 2 from burning fossil fuel. Powdered limestone decomposes into CaO, which reacts with S0 2 to form CaS0 3' The remaining S0 2 is combined with an aqueous suspension of CaO to form CaS0 3' 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, whereas 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. All in all, some 50 million to 60 million tons of S02 are released into the atmosphere each year! In the troposphere, SO? is almost all oxidized to H 2 S0 4 in the form of aerosol, which ends up in wet precipitation or acid rain. The mechanism for the conversion of S02 to H 2 S0 4 is quite complex and not fully understood. The reaction is believed to be initiated by the hydroxyl radi- cal (OH): • The HOSO? radical is further oxidized to S0 3: The sulfur trioxide formed would then rapidly react with water to form sulfuric acid: S0 2 can also be oxidiz ed to S0 3 and then converted to H 2 S0 4 on particles by heterogeneous catalysis. Eventually, the acid rain can corrode limestone and marble (CaC0 3 ). A typical reac- tion is Sulfur dioxide can also attack calcium carbonate directly: There are two ways to minimize the effects of S02 pollution. The most direct approach is to remove sulfur from fossil fuels before combustion, but this is technologically difficult to accom- plish. A cheaper but less efficient way is to remove S0 2 as it is formed. For example, in one process powdered limestone is injected into the power plant boiler or furnace along with the coal (Figure 21.23). At high temperatures, the following decomposition occurs: CaC0 3 (s ) CaO(s) + CO 2 (g) limestone quicklime The quicklime reacts with S0 2 to form calcium sulfite and some calcium sulfate: CaO(s) + S02(g) CaS03(S) 2CaO(s) + 2S0ig) + 0 2(g) • 2CaS04(s) S + 0 2 ' S0 2 CaC0 3 • CaO + CO 2 CaO + S0 2 • CaS0 3 Furnace Coal Mostly CO 2 and air ,. Smokestack Purification chamber nnll Aqueous suspension of CaO SECTION 21.7 Photochemical Smog 845 To remove any remaining S02, 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 caJled liming to reduce their acidity (Figure 21.24). Installing a sulfuric acid plant near a metal ore refining site is also an effective way to cut SOz emission because the S0 2 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! 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, photochemical smog, which is formed by the reactions of automobile exhaust in the presence of sunlight, is much more common. 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 N0 2 and 0 3 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: Once released into the atmosphere, nitric oxide is oxidized to nitrogen dioxide: Sunlight causes the photochemical decomposition of N0 2 (at a wavelength shorter than 400 nm) into NO and 0: NOig) + hv +. 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: where M is some inert substance such as N 2 . Ozone attacks the C=C linkage in rubber: R R \ / C=C + 03 / \ R R • R R \ /0., / C/ 'c 1\ j\ ROO R H2 0 • R \ C=O / 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 hydro- carbons, nitrogen oxides, and oxygen. One of the products of these reactions is peroxyacetyl nitrate (PAN): CH3-C-0-0-N02 II o PAN is a powerful lachrymator, or tear producer, and causes breathing difficulties. Figure 21.25 shows typical variations with time of primary and secondary pollutants. Ini- tially, the concentration of N0 2 is quite low. As soon as solar radiation penetrates the atmosphere, though, more N0 2 is formed from NO and O 2 . The concentration of ozone remains fairly constant at a low level in the early morning hours. As the concentration of unburned hydrocarbons and alde- hydes increases in the air, the concentrations of N0 2 and 0 3 also rise rapidly. The actual amounts depend on the location, traffic, and weather conditions, but their presence is always accompanied by haze (Figure 21.26). The oxidation of hydrocarbons produces'var~ous organic intermediates, such as alcohols and carboxylic acids, which are all less vo lati~ than the hydrocarbons them- selves. These substances eventually condense into small droplets of llq~. The dispersion of these droplets in air, called an aerosol, scatters sunlight and reduces visibility. This interaction also makes the air look hazy. Figure 21 .24 Spreading calcium oxide (CaO) over acidified soil. This process is called limjng. • 846 CHAPTER 21 Environmental Chemistry Figure 21 .25 Typical variations with time in concentration of air pollutants on a s moggy day. • Figure 21.26 A smoggy day in a big city. NO;.2 _ NO 4 6 8 10 A. M. Hydrocarbons 12 Noon 2 4 6 P.M. 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 CO 2 and H 2 0 and to reduce NO and N0 2 to N2 and O 2 [ ~~ Section 14.6]. 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: 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" warn- ings frequently issued to its residents. Indoor Pollution Difficult as it is to avoid air pollution outdoors, it is no easier to avoid pollution indoors. 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, car- bon monoxide, carbon dioxide, and formaldehyde. The Risk from Radon In a highly publicized case about 35 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 con- tamination 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? Radon is a member of Group 8A (the noble gases). It is an intermediate product of the radioactive decay of uranium-238 (see Figure 20.3). 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 21.27). 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 build- ing through tiny cracks in the ba sement floor. It is slightly soluble in water, so it can be spread in different media. Radon-222 is an a-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 a particles emitted by polonium and its decay products, which are also radioactive, can cause lung cancer. SECTION 21.8 Indoor Pollution 847 Other human-made radiation 3% Radon 54% Nuclear ~~~?/~ : I testing Other natural 2% radiation Medical procedures 14 % 27% 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 21.28). 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 (I' 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 Q' particle will leave a "track." After several months' exposure, the film is etched with a sodium hydroxide solution and the number 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 (e.g., 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 smok- ers. It seems quite likely that there is a synergistic effect between radon and smoking on the development of lung cancer. Radon decay 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 COf- responds to 3.70 X 10 10 disintegrations of radioactive nuclei per second; a picocurie is a trillionth of a curie, Of 3.70 X 10- 2 disintegrations per second.) • :Sample Problem 21.3 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? Strategy All radioactive decays obey first-order kinetics, making the half-life independent of the initial concentration. Setup Because the question involves an integral number of half-lives, we can deduce the amount of Rn-222 remaining without using Equation 14.3. Solution After one half-life, the amount of Rn left is 0.5 X 1.0 g, or 0.5 g. After two half-lives, only 0.25 g of Rn remains. Generalizing the fraction of the isotope left after n half-lives as (1/2)", where n = 10, we write quantity of Rn-222 left = 1.0 g X (1/2)10 = 9.8 X 10- 4 g Practice Problem The concentration of Rn-222 in the basement of a house is 1.8 X 10- 6 mol/L. Assume the air remains static, and calculate the concentration of the radon after 2.4 days. Figure 21.27 Sources of background radiation. Fig u re 21.28 Home radon detector. Think About It An alternative solution is to calculate the first- order rate constant from the half- life and use Equation 14.3: N In t = -kt No where N is the mass of Rn-222. Try this nd verify that your answers are the same. (Since most of the kinetics problems we encounter do not involve an integral number of half-lives, we generally use Equation 14.3 as the first approach to solving them.) [...]... argued, based in part on the work of Paul Crutzen, that to do so would significantly endanger the ozone layer Problems: a) The bond enthalpy of NO is 630.6 kJ/mol Determine the maximum wavelength of light required to break the bond in an NO molecule [ ~~ Sample Problem 21.1] b) Can nitric oxide act as a greenhouse gas? Explain [ ~~ Sample Problem 21.2] • 849 850 CHAPTER 21 Environmental Chemistry CHAPTER... 832 UESTIONS AND PROBLEMS • Section 21.1: Earth's Atmosphere Problems Review Questions 21.5 Referring to Table 21.1, calculate the mole fraction of CO 2 and its concentration in parts per million by volume 21.6 Calculate the partial pressure of CO2 (in atm) in dry air when the atmospheric pressure is 754 mmHg 21.7 Describe the processes that result in the warming of the stratosphere 21.8 Calculate the... (c) Under the conditions described in part (b), the half-life of the reaction has been estimated to be 6.4 X 103 min What would the half-life be if the initial concentration of NO were 10 ppm? Section 21.6: Acid Rain 21.43 Assume that the formation of nitrogen dioxide: 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... (DMSA), marketed under the name Chemet Chelation therapy relies on coordination chemistry o Covers And Lasts longer 1 Than "OLD DUTCH " Process While Le:.~ J SH OH HO SH 0 Dimercaptosuccinic acid (DMSA) In This Chapter You, Will Learn about the properties of coordination compounds and how, through the use of chelates, coordination chemistry is used to solve a variety of medical and other societal problems... 54 I Xe 85 86 At Rn (ll7) 118 1 Alfred Werner (1866- 1919) Swiss chemist Werner started as an organic chemist but did his most notable work in coordination chemistry For his theory of coordination compounds, Werner was awarded the Nobel Prize in Chemistry in 1913 SECTION 22.1 Sc Ti Coordination Compounds 859 The trends in radii, electronegativity, and ionization energy for the first row of transition... Identify the components of each compound, and use known oxidation states and charges to determine the oxidation state of the metal ( Continued) 862 CHAPTER 22 Coordination Chemistry Setup (a) [Ru(NH3)s(H2 0)]CI 2 consists of a complex ion (the part of the formula enclosed in square brackets) and two Cl- counter ions Because the overall charge on the compound is zero, the complex ion is [Ru(NH3)s(H2 0)f + There... Based on this information, explain why the C-CI bond in a CFC molecule is preferentially broken by solar radiation at 250 nm 21.26 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: 21.9 What process gives rise to the aurora borealis and aurora australis? 21.10 Why can astronauts not release... dioxide? Problems 21.39 • The annual production of zinc sulfide (ZnS) is 4.0 X 104 tons Estimate the number of tons of S02 produced by roasting it to extract zinc metal 852 21.40 CHAPTER 21 Environmental Chemistry Calcium oxide or quicklime (CaO) is used in steelmaking, cement manufacture, and pollution control It is prepared by the thermal decomposition of calcium carbonate: CaC0 3(s) - _ CaO(s) + CO... b) 0.0040 g d) coma c) 0.99 g e) lung cancer d) 0.97 g e) 0.0010 g • APPLYING WHAT YOU'VE LEARNED Applying What You've Learned The research of Paul Crutzen, the third recipient of the Nobel Prize for Chemistry in 1995, involved the effect of nitric oxide (NO) on the destruction of stratospheric ozone Unlike CFCs, which may take 50 to 100 years to diffuse into the upper atmosphere, nitric oxide is introduced... S02(g) + H 20 (l) • H +(aq) + HS0 3"(aq) Given that the equilibrium constant for the preceding reaction is 2 1.3 X 10- , calculate the pH of the rainwater Assume that the reaction does not affect the partial pressure of S02' Section 21.7: Photochemical Smog 21.61 List the major indoor pollutants and their sources 21.62 What is the best way to deal with indoor pollution? 21.63 Why is it dangerous to . photos were taken just a few decades apart. ,;' Multimedia Organic and Biochemistry-oil ref i ning processes. 844 CHAPTER 21 Environmental Chemistry Figure 21.22 Sulfur dioxide. is toward an increase in CO 2 , The current rate of increase is about 1 ppm (1 part CO 2 per million parts air) by volume per year, which is equivalent to 9 X 10 9 tons of CO 2 ! Scientists. 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