The atmosphere
Even though large quantities of iron coalesced
into Earth's core, appreciable quantities remained in the mantle and crust. Today's crust, for example, is about 5 percent iron by weight. The iron that remained near the surface during early Earth played an important role in concentrating oxygen in Earth's early atmosphere. Oxygen gas dissolves in water. As a result, when early cyanobacteria released oxygen, some of the gas dissolved in seawater. In an oxygen-poor environment, iron also dissolves in water. However, when oxygen is abundant, the dissolved oxygen reacts with dissolved iron and causes the iron to precipitate rapidly as iron-oxide minerals. Because iron is soluble in water lacking dissolved oxygen, large amounts of iron dissolved in the early Precambrian Ocean at times when its concentration of dissolved oxygen was below the threshold level. However, when the concentration of dissolved oxygen from cyanobacteria rose beyond the threshold, iron that had been dissolved in the seawater precipitated rapidly as iron-oxide minerals. These minerals settled to the seafloor, forming a layer there (Figure 17.3). This process also removed oxygen from the water, explaining why the oxygen concentration in the atmosphere didn't rise at this time in Earth's early history, even though cyanobacteria were releasing oxygen as a gas. When the oxygen concentration in the atmosphere and seawater reached a threshold, the oxygen combined with dissolved iron to form iron oxide minerals. The iron oxide minerals precipitated from the water and settled to the seafloor, forming the iron-rich layer of a banded iron formation. Roughly 90 percent of the iron ore that is mined globally comes from banded iron formation, chemical sedimentary rock that consists of alternating layers ("bands") of iron oxide and chert a few centimeters thick (Figure 17.4). Most of the banded iron formation on Earth is between 2.6 and 1.9 billion years old and is inferred to have originated when the oxygen level in the seas hovered near the threshold at which soluble iron combined with oxygen to form iron-oxide minerals. At times when dissolved oxygen concentrations were high enough, the oxygen would combine with dissolved iron to precipitate iron-oxide minerals, which formed a layer on the seafloor. The formation of these minerals extracted oxygen as well as iron from the seawater, lowering its oxygen concentration below the threshold. Then dissolved iron accumulated again in the seas from chemical weathering of exposed rock, while the oxygen was slowly replenished by photosynthesizing cyanobacteria. During that time, clay and other minerals washed from the continents and accumulated on the seafloor as they do today, forming the thin layers of silicate minerals that lie between the iron-rich layers in banded iron formation. When the oxygen concentration rose above the threshold again, another layer of iron oxide minerals formed. In this banded iron formation from Michigan, the red bands are iron oxide minerals and the dark layers are chert (silica). Banded iron formations contain thousands of alternating layers of iron minerals and silicates and can cover tens of square kilometers. The great volume of the iron formations, coupled with the fact that they continued to form from 2.6 to 1.9 billion years ago, suggests that the reactions that formed them must have kept the levels of dissolved oxygen close to the threshold for about 700 million years. Thus the iron-rich rocks that support our industrial society were formed by interactions among early photosynthesizing organisms, sunlight, air, and the oceans. In the primordial atmosphere, free oxygen also reacted with hydrogen to form water. This process helped keep the oxygen concentration in the atmosphere low. However, after life evolved, bacteria in the oceans removed atmospheric hydrogen in a process that produced methane. When the hydrogen concentration decreased sufficiently, the concentration of free oxygen in the atmosphere could rise again. Whatever the exact combination of biological and geochemical processes, evidence in the rocks indicates that the oxygen concentration in the atmosphere remained low and then jumped suddenly from trace to appreciable quantities approximately 2.4 billion years ago.
Coal is
largely carbon, which, when burned completely, produces carbon dioxide. Petroleum is a mixture of hydrocarbons, compounds composed of carbon and hydrogen. When hydrocarbons burn completely, they produce carbon dioxide and water as the only combustion products. Neither is poisonous, but both are greenhouse gases. If fuels were composed purely of compounds of carbon and hydrogen, and if they always burned completely, air pollution from the burning of fossil fuels would pose little direct threat to our health (although combustion of fossil fuels would still contribute to global warming). However, fossil fuels contain impurities, and combustion is usually incomplete. As a result, other products form—most of which are harmful. Products of incomplete combustion include hydrocarbons such as benzene and methane. Benzene is a carcinogen (a compound that causes cancer), and methane is another greenhouse gas. Incomplete combustion of fossil fuels releases many other pollutants, including carbon monoxide (CO), which is colorless and odorless yet very toxic. Additional problems arise because coal and petroleum contain impurities that generate other kinds of pollution when they are burned. Small amounts of sulfur are present in coal and, to a lesser extent, in petroleum. When these fuels burn, the sulfur forms oxides, mainly sulfur dioxide SO2 and sulfur trioxide SO3 . High sulfur dioxide concentrations have been associated with major air pollution disasters of the type that occurred in Donora. Today, the primary global source of sulfur dioxide pollution is electricity generation by coal-fired power plants. Nitrogen, like sulfur, is common in living tissue and therefore is found in all fossil fuels. This nitrogen, together with a small amount of atmospheric nitrogen, reacts when coal or petroleum is burned. The products are mostly nitrogen oxide (NO) and nitrogen dioxide NO2 . Nitrogen dioxide is a reddish-brown gas with a strong odor. It contributes to the "browning" and odor of some polluted urban atmospheres. Automobile exhaust is the primary source of nitrogen oxide pollution.
The carbon dioxide
water, and nitrogen atmosphere changed as volatiles escaped from Earth's mantle to the surface in volcanic eruptions, in a process called outgassing. In 1953, Stanley Miller and Harold Urey hypothesized—on the basis of little direct modeling—that by 4 billion years ago, Earth's atmosphere consisted primarily of methane CH4 , ammonia NH3 , hydrogen H2 , and water H2O . They mixed these gases in a glass flask and fired sparks across the flask to simulate Earth's early atmosphere, beset by lightning storms. Amino acids, the building blocks of proteins, formed in the flask. The Miller-Urey model immediately became popular because scientists speculated that the first living organisms formed by accretion of these abiotic (nonliving) amino acids. In the early 1970s, the idea of a methane-ammonia-hydrogen atmosphere was largely discredited, and most scientists postulated that the Hadean atmosphere was composed mainly of carbon dioxide CO2 , with smaller amounts of nitrogen N2 , water H2O , and other gases. However, in 2002, geologists analyzed the isotopic geochemistry of single grains of zircon that are the oldest known preserved mineral grains on Earth, dating between 4.5 and 4.0 billion years before present. Zircon is a silicate mineral that only forms in rocks with a silica-rich, felsic composition; it does not exist in the silica-poor mafic and ultramafic rocks, so the geologists concluded that granitic crust had formed within the first 500 million years of Earth's history. In addition, the isotopic composition of oxygen making up the silicate anionic groups within the zircon mineral crystal was high enough for the geologists to conclude that the surface of the early Earth was cold, not hot, and that water would have been mostly liquid. Subsequently, in 2005, scientists used models of Earth's Hadean geosphere to calculate how gases trapped in the interior would react with rocks and minerals in the planet's interior and surface. These results also suggested that early Earth had a cool surface that included liquid water. From this work, one unobvious, yet pivotal, question emerged: When did Earth's core form? Earth's core is the deepest layer of the geosphere. The atmosphere is a thin veneer surrounding the crust. How could one affect the other? Recall that the core is composed primarily of iron and nickel. Before Earth's interior segregated into a layered core and mantle, the mantle contained more iron than it does today. Then, when iron settled into the core, the mantle became relatively iron-poor. Modern hypotheses state that Earth was hot at the time of its formation and that most of the planet's iron was sequestered in its core shortly after the planet evolved, leaving an iron-poor mantle behind. Because mantle rocks rise to the surface through volcanic eruptions, the composition of the mantle directly influences the composition of erupted material, including the types of volatiles that are emitted. By modeling the reactions of volatiles in an iron-poor mantle, researchers concluded in 2005 that the delivery of these volatiles to the atmosphere resulted in a composition that included not only large amounts of hydrogen, as Miller and Urey postulated, but also high concentrations of carbon dioxide, as later models had proposed. Even more recently, a different set of scientists argued that the Earth's early surface temperatures were significantly warmer than today, with estimates of surface temperatures between 45*C and 85*C. Models of Earth and its evolution are constantly changing. Scientists propose hypotheses and theories, consider new ideas—and often other scientists disagree. This is the nature and the joy of science. In a review article, Christopher Chyba of SETI Institute and Stanford University wrote that current disagreement about the composition of Earth's earliest atmospheres "makes it a great time for young scientists to enter the field, but it also reminds us that some humility regarding our favorite models is in order." Further constraining the composition and evolution of Earth's earliest atmosphere continues to be a major scientific challenge that will involve much additional research, discussion, and—quite likely—disagreement among different scientists.
Imagine that
your great-grandfather had entered the exciting new business of making moving pictures. Old-time photographic film was "slow" and required lots of sunlight, so many filmmakers left the polluted, overcast, industrial Northeast. Southern California, with its warm, sunny climate and little need for coal, was preferable. Thus, a district of Los Angeles called Hollywood became the center of the movie industry. Its population boomed, and after World War II, automobiles became about as numerous as people. Then the quality of the air deteriorated in a strange way. People noted four different kinds of changes: (1) a brownish haze called smog settled over the city (Figure 17.16); (2) people felt irritation in their eyes and throats; (3) vegetable crops became damaged; and (4) the sidewalls of rubber tires developed cracks. In the 1950s, air pollution experts worked mostly in the industrialized cities of the East Coast and the Midwest. When they were called to diagnose the problem in Southern California, they looked for the sources of air pollution they knew well, especially sulfur dioxide. But the smog was nothing like the pollution they were familiar with. These researchers eventually learned that incompletely burned gasoline in automobile exhaust reacts with nitrogen oxides and atmospheric oxygen in the presence of sunlight to form ozone O3 . The ozone then reacts further with automobile exhaust to form smog (Figure 17.17). Smog forms in a sequential process. Step 1: Automobile exhaust reacts with air in the presence of sunlight to form ozone. Step 2: Ozone reacts with automobile exhaust to form smog. As you learned in section on "Atmospheric Temperature," ozone in the stratosphere absorbs ultraviolet radiation and protects life on Earth. Yet, excessive ozone in the air we breathe is harmful. Is ozone a pollutant to be eliminated, or a beneficial component of the atmosphere that we want to preserve? The answer is that it is both, depending on where it is found: ozone in the troposphere reacts with automobile exhaust to produce smog, and therefore it is a pollutant. Ozone in the stratosphere is beneficial, and the destruction of the ozone layer there creates serious problems. Ozone irritates the respiratory system, causing loss of lung function and aggravating asthma in susceptible individuals. Ozone also increases susceptibility to heart disease and is a suspected carcinogen. High ozone concentrations slow the growth of plants, which is a particularly serious problem in the rich, agricultural areas of California. Ozone pollution ranks as the 33rd leading cause of human death. As of 2018, over 95 percent of Earth's human population lives with unsafe levels of air pollution, using air quality guidelines from the World Health Organization. Using population-adjusted measurements of ozone, with greater weight given to ozone concentrated in densely populated areas, the State of Global Air/2018 reported that global levels of ozone pollution increased between 1990 and 2016. Within the United States, ozone levels fell during the same period, as air pollution controls and fuel emission standards were implemented.
aerosol
In pollution terminology, a particle or particulate that is suspended in air.
particle, or particulate
In pollution terminology, any small piece of solid matter larger than a molecule, such as dust or soot.
chlorofluorocarbons (CFCs)
Organic compounds containing chlorine and fluorine, which rise into the upper atmosphere and destroy the ozone layer there.
Atmospheric pressure, often called barometric pressure
The pressure of the atmosphere at any given location and time.
tropopause
The top of the troposphere; the boundary between the troposphere and the stratosphere.
bar
Unit of measurement for atmospheric pressure. One bar is roughly equal to atmospheric pressure at sea level.
smog
Visible, brownish air pollution formed through chemical reactions that involve incompletely combusted automobile gasoline, nitrogen dioxide, oxygen, ozone, and sunlight.
barometer
A device used to measure barometric pressure.
pH scale
A logarithmic scale that measures the acidity of a solution. A pH of 7 is neutral; numbers lower than 7 represent acidic solutions, and numbers higher than 7 represent basic ones.
banded iron formation
A marine chemical sedimentary rock formed mostly between 2.7 and 1.9 billion years ago and consisting of centimeter-scale interbeds of iron oxide and chert. Formed as atmospheric oxygen alternately went through periods of accumulation as a waste product from photosynthesizing cyanobacteria and periods of withdrawal during widespread precipitation of iron oxide minerals.
cyanobacteria
Blue-green algae that were among the earliest photosynthetic life-forms on Earth.
thermosphere
An extremely high and diffuse region of the atmosphere lying above the mesosphere, from about 80 kilometers upward.
ozone hole
An unusually low ozone concentration in the stratosphere that is centered roughly over Antarctica.
halons
Compounds containing bromine and chlorine, which rise into the upper atmosphere and destroy the ozone layer there.
fly ash
Noncombustible minerals that escape into the atmosphere when coal burns, eventually settling as gritty dust.
acid precipitation, also called acid rain
Rain, snow, fog, or mist that has become acidic after reacting with air pollutants.
stratopause
The ceiling of the stratosphere; the boundary between the stratosphere and the mesosphere.
mesopause
The ceiling of the mesosphere; the boundary between the mesosphere and the thermosphere.
stratosphere
The layer of air above the tropopause, extending upward to about 55 kilometers.
mesosphere
The layer of air that lies above the stratopause, extending upward from about 55 kilometers to about 80 kilometers above Earth's surface.
troposphere
The layer of air that lies closest to Earth's surface and extends upward to about 17 kilometers.
exosphere
The outermost layer of the in which gas molecules are so diffuse they do not collide. Merges outward into interstellar space.
photosynthesis
The process by which chlorophyll-bearing plant cells convert carbon dioxide and water to organic sugars, using sunlight as an energy source; oxygen is released in the process.
Second Great Oxidation Event
The process, occurring about 600 million years ago, when the oxygen concentration in Earth's atmosphere increased abruptly a second time in response to interactions among chemical, physical, and biological processes.
outgassing
The release of volatiles from Earth's mantle and crust during volcanic eruptions at the surface
First Great Oxidation Event
The sudden increase in Earth's atmospheric oxygen concentration from trace amounts to appreciable quantities that occurred approximately 2.4 billion years ago, probably because of a combination of biological and geochemical processes.
Gaia
The term (Greek for "Earth") used by James Lovelock to refer to our planet, which he likened to a living creature due to the interconnectivity of all of Earth's systems.
The molecules in a gas zoom
about in a random manner. For example, at 20*C an average oxygen molecule is traveling at 425 meters per second (950 miles per hour). In the absence of gravity, or where there are temperature differences or other perturbations, a gas will fill a space homogeneously. Thus, if you float a cylinder of gas in space, the gas would disperse until there is an equal density of molecules and equal pressure throughout the cylinder. In contrast, gases that surround Earth are perturbed by many influences, which ultimately create the complex and turbulent atmosphere that helps shape the world we live in. Within our atmosphere, gas molecules zoom around, as in the imaginary cylinder, but in addition, gravity pulls them downward. As a result of this downward force, more molecules concentrate near the surface of Earth than at higher elevations. Therefore, the atmosphere is denser at sea level than it is at higher elevations—and the pressure is higher. Density and pressure then decrease exponentially with elevation (Figure 17.6). At an elevation of about 5,000 meters, the atmosphere contains about half as much oxygen as it does at sea level. If you ascended in a balloon to 16,000 meters (16 kilometers) above sea level, you would be above 90 percent of the atmosphere and would need an oxygen mask to survive. At an elevation of 100 kilometers, pressure is only 0.00003 that of sea level, approaching the vacuum of outer space. There is no absolute upper boundary to the atmosphere. Atmospheric pressure decreases with altitude. One-half of the atmosphere lies below an altitude of 5,600 meters. Atmospheric pressure, often called barometric pressure, is measured with a barometer. A simple but accurate barometer is constructed from a glass tube that is sealed at one end. The tube is evacuated and the unsealed end placed in a dish of a liquid such as mercury. The mercury rises in the tube because atmospheric pressure depresses the level of mercury in the dish, but there is no air in the tube to counter that pressure so the surface of the mercury rises (Figure 17.7). At sea level mercury rises approximately 76 centimeters, or 760 millimeters (about 30 inches), into an evacuated tube. (A) Atmospheric pressure forces mercury upward in an evacuated glass tube. The height of the mercury in the tube is the measure of air pressure. (B) Three common scales for reporting atmospheric pressure and the conversion among them. Meteorologists express pressure in inches or millimeters of mercury, referring to the height of the column of mercury in a barometer. They also express pressure in bars and millibars. A bar is approximately equal to sea level atmospheric pressure. A millibar is 0.001 of a bar. A mercury barometer is a cumbersome device nearly a meter tall, and mercury vapor is poisonous. A safer and more portable instrument for measuring pressure, called an aneroid barometer, consists of a partially evacuated metal chamber connected to a pointer. When atmospheric pressure increases, it compresses the chamber and the pointer moves in one direction. When pressure decreases, the chamber expands, directing the pointer the other way (Figure 17.8). In an aneroid barometer, increasing air pressure compresses the air-tight chamber and causes the connected pointer to move in one direction. When the pressure decreases, the chamber expands, deflecting the pointer the other way. Changing weather can also affect barometric pressure. On a stormy day at sea level, pressure may drop to 980 millibars (28.94 inches), although barometric pressures below 900 millibars (26.58 inches) have been reported during some hurricanes. In contrast, during a period of clear, dry weather, a typical high-pressure reading may be 1,025 millibars (30.27 inches). These changes are discussed more in Chapter 19.
Several deposits of banded iron formed
after the First Great Oxidation Event, and this process continued to remove oxygen that was released during photosynthesis. The last, major banded iron layer was deposited about 1.9 billion years ago, but the oxygen concentration in the atmosphere did not increase dramatically. At least two more critical steps were required before efficient multicellular organisms could evolve. The Sun emits energy largely in the form of high-energy ultraviolet light. These rays are energetic enough to break apart complex molecules and kill evolving multicellular organisms. But high-altitude oxygen absorbs ultraviolet radiation in a process that forms ozone O3. Thus, the oxygen concentration couldn't increase in the lower atmosphere until appreciable concentrations first accumulated in the upper atmosphere. To summarize: oxygen, largely produced by the earliest photosynthetic organisms, was not only necessary for life as we know it today but as ozone it also protected multicellular life by filtering out harmful solar rays. Multicellular plants and animals emerged in Late Precambrian time, between 1 billion and 543 million years ago. About 600 million years ago, the oxygen level in the atmosphere increased rapidly a second time, in a process called the Second Great Oxidation Event. What changed abruptly 1.3 billion years after the last banded iron layers were deposited to allow oxygen to accumulate? Scientists propose that prior to 600 million years ago, biological decay was almost as rapid as photosynthesis. Therefore, the oxygen that was released into the atmosphere was immediately consumed during respiration, according to the following reactions: During Photosynthesis: Carbon Dioxide + Water turns into sugars + Oxygen. During respiration and decay: Sugars + Oxygen turns into Carbon Dioxide + Water. Then, beginning abruptly 600 million years ago, several processes occurred that allowed preservation of organic matter in sediments before it could decay. These processes effectively removed the organic carbon (the sugars) from the system, thereby causing the opposing reactions above to lean in the direction favoring photosynthesis. Stated differently, environmental conditions that allowed the generation but subsequent removal and storage of organic carbon in sediment also were conducive to the expansion of photosynthetic organisms and the generation of much oxygen as a by-product. All the proposed processes for sequestering organic matter involve complex reactions among Earth's four spheres: Geochemical processes produced an abundance of clays 600 million years ago. These clays buried and preserved organic matter on the seafloor. Zooplankton evolved in the seas. These organisms produced dense, organic-laden feces that sank rapidly to the seafloor, where they accumulated in the clays (mentioned above) and were subsequently buried, thereby removing organic carbon from the system. Simple lichens evolved on land. The lichens accelerated weathering, and the weathered ions washed into the sea and provided nutrients for phytoplankton. In turn, the phytoplankton fed the zooplankton, which sequestered nutrients as described previously. Thus, numerous complex chemical, physical, and biological processes combined to set the stage for the Second Great Oxidation Event. Earth's atmosphere not only sustains us but it insulates Earth's surface as winds distribute the Sun's heat around the globe so that the surface is neither too hot nor too cold for life to exist. Clouds form from water vapor in the atmosphere and rain falls from clouds. In addition, the atmosphere filters out much of the Sun's ultraviolet radiation, which can destroy living tissue and cause cancer. The atmosphere carries sound; without air we would live in silence. Without an atmosphere, airplanes and birds could not fly; wind would not transport pollen and seeds; the sky would be black rather than blue; and no reds, purples, and pinks would color the sunset. When we understand this complex web of interacting processes and realize that we are the only planet in the Solar System to be so fortunate, we can only marvel at the fragility of Earth's atmosphere.
The temperature of the atmosphere changes with
altitude (Figure 17.9). The layer of air closest to Earth—the layer we live in—is the troposphere. Virtually all of the water vapor and clouds exist in this layer, and almost all weather occurs here. Earth's surface absorbs solar energy, and thus the surface of the planet is warm. But, as explained earlier, continents and oceans also radiate heat, and some of this energy is absorbed by the troposphere. Lower parts of the troposphere absorb most of the heat radiating from Earth's surface; in contrast, at higher elevations in the troposphere, the atmosphere is thinner and absorbs less energy. Consequently, temperature decreases with increased elevation in the troposphere; mountaintops are generally colder than valley floors, and commercial jet airliners must heat their cabins once reaching cruising altitudes, which are generally between 10 and 11 kilometers elevation. Atmospheric temperature varies with altitude. The atmospheric layers are zones in which different factors control the temperature. The top part of the troposphere is the tropopause, which lies at an altitude of about 17 kilometers at the equator, although it is lower at the poles. The tropopause is the boundary between the troposphere and the stratosphere above. It is characterized by an abrupt cessation in the steady decline in temperature with altitude, because cold air from the upper troposphere is too dense to rise higher. As a result, little mixing of air molecules occurs across the tropopause. The tropopause forms the floor of the stratosphere, in which temperature remains constant to 35 kilometers and then increases with altitude until, at about 50 kilometers, it is as warm as air at the Earth's surface. Although the stratosphere is above elevations at which commercial air liners can fly, many military jets can reach it if they really try; the stratosphere is not too high to become polluted by radioactive elements from atmospheric tests of atomic and hydrogen bombs. Large levels of anthropogenic radionuclides (specific isotopes of radioactive elements known only to come from human-caused nuclear reactions) were released into the atmosphere during atmospheric tests of nuclear weapons conducted mostly in the 1950s and 1960s by the United States and former Soviet Union and in the 1970s by France and China. Most of the radioactive debris from atmospheric test of hydrogen bombs reached the stratosphere and accumulated there, forming a reservoir of radioactive radionuclides. For many years, global fallout of these radionuclides from the stratosphere downward through the troposphere to Earth's surface was observed during the late spring when heating of the ground surface and lower troposphere formed rising hot air that broke through the tropopause and disrupted the lower stratosphere. This disruption caused cold, radionuclide-rich air from the lower stratosphere to rapidly descend, causing the nuclear fallout. Each spring beginning prior to the 1963 moratorium on atmospheric testing of nuclear weapon until the early 1990s, the annual rate of radionuclide fallout peaked, although the annual peak concentration fell through this period. Since the 1990s, the biggest fluctuations in the concentration of radionuclides in the troposphere are due mostly to their resuspension from contaminated soil (Figure 17.10). Thus, although the annual springtime "radionuclide fall-out event" no longer is significant in terms of providing the source of radionuclides in the troposphere, the deposition of radioactive elements during those events and the physical reworking of those deposits today provide the main source of atmospheric radionuclides in the troposphere. (A) Measured levels of Plutonium-239 and Plutonium-240 concentrations in the troposphere between 1960 and 2010. Following the moratorium on atmospheric testing of nuclear weapons in 1963, the concentration of the two radionuclides dropped in the troposphere and the stratosphere. (B) Atmospheric nuclear weapons tests by the Chinese during the 1970s and early 1980s led to high sustained and high concentrations of Plutonium 239 and Plutonium-240 in the stratosphere. Following the elimination of these tests, stratospheric concentrations of the Plutonium radionuclides dropped. Katsumi Hirose and Pavel P. Povinec, "Sources of plutonium in the atmosphere and stratosphere-troposphere mixing", www.nature.com/scientificreports, 28 October 2015. The reversal in the temperature profile between the troposphere and the stratosphere occurs because the two atmospheric layers are heated by different mechanisms. As already explained, the troposphere is heated primarily from below, by Earth. The stratosphere, however, is heated primarily from above, by direct incoming solar radiation. Oxygen molecules O2 in the stratosphere absorb energetic ultraviolet rays from the Sun. The radiant energy breaks the oxygen molecules apart, releasing free oxygen atoms. These free oxygen atoms then recombine to form ozone O3 . Ozone absorbs ultraviolet energy more efficiently than oxygen does, and the absorption of UV radiation by the ozone warms the upper stratosphere. Ultraviolet radiation is energetic enough to affect organisms. Small quantities give us a suntan, but large doses cause skin cancer and cataracts of the eye, inhibit the growth of many plants, and otherwise harm living tissue. The ozone in the upper atmosphere protects life on Earth by absorbing much of this high-energy radiation before it reaches Earth's surface. Ozone concentration declines in the upper portion of the stratosphere, and therefore at about 55 kilometers above Earth, temperature once more begins to fall rapidly with elevation. This boundary between rising and falling temperature is the stratopause, the ceiling of the stratosphere. The second zone of declining temperature in Earth's modern atmosphere is the mesosphere. Little radiation is absorbed in the mesosphere, and the thin air is extremely cold. The ceiling of the mesosphere is the mesopause. Starting at about 80 kilometers above Earth, the temperature again remains constant and then rises rapidly in the thermosphere. Here the atmosphere absorbs high-energy X-rays and ultraviolet radiation from the Sun. High-energy reactions strip electrons from atoms and molecules, producing ions. The temperature in the upper portion of the thermosphere is just below freezing—not extremely cold by surface standards. The uppermost layer of the atmosphere is the exosphere a layer of very diffuse gas that overlies the thermosphere and thins upward into the vacuum of space. In the exosphere, gas molecules are gravitationally influenced by to the Earth, but the density of gas is too low for the individual molecules to collide with each other, as gas normally does at lower atmospheric levels. Instead of colliding with other molecules, gas molecules in the exosphere travel along 'ballistic trajectories' similar to the arc of a thrown ball. As they are pulled by Earth's gravity, slower gas molecules arc back into the thermosphere while some of the faster-moving molecules continue into outer space and are lost.
chemoautotrophs
an organism, typically a bacterium, that derives its metabolic energy by oxidizing inorganic compounds.
A particle, or particulate, is
any small piece of solid matter, such as dust or soot. An aerosol is any small particle that is larger than a molecule and suspended in air. In the context of air pollution, all three terms are used interchangeably. Many natural processes release aerosols. Windblown silt, pollen, volcanic ash, salt spray from the oceans, and smoke and soot from wildfires are all aerosols. Industrial emissions add to these natural sources. Smoke and soot are carcinogenic aerosols formed whenever fuels are burned. Coal always contains clay and other noncombustible minerals that have accumulated along with the organic matter in the depositional environment in which the coal formed, commonly a swamp. When the coal burns, some of these minerals escape from the chimney as fly ash, which settles as gritty dust. When metals are mined, the drilling, blasting, and digging raise dust, and this too adds to the total load of aerosols. In 1988, EPA epidemiologists noted that whenever atmospheric aerosol levels rose above a critical level in Steubenville, Ohio, the number of fatalities from all causes—car accidents to heart attacks—rose. After several studies substantiated the Steubenville report, the EPA proposed additional reductions of ambient aerosol levels in the United States. Opponents argued that it is unfair to target all aerosols, because the term covers a wide range of substances from a benign grain of salt to a deadly mist of toxic volatiles.
Our Solar System formed
from a cold, diffuse cloud of interstellar gas and dust. About 99.8 percent of this cloud was composed of the two lightest elements: hydrogen and helium. Consequently, when Earth formed, its primordial atmosphere was composed almost entirely of these two light elements. But because Earth is relatively close to the Sun and its gravitational force is relatively weak, its primordial hydrogen and helium atmosphere rapidly boiled off into space and escaped. Table 17.1 shows this and the subsequent atmospheres of Earth described in this chapter. In Chapter 15, we learned that most of the volatile compounds that form Earth's hydrosphere, atmosphere, and biosphere originated from outer parts of the Solar System. Recall that, in its infancy, the Solar System was crowded with bits of rock, comets, ice chunks, and other debris left over from the initial coalescence of the planets. These bolides crashed into the planet in a near-continuous rain that lasted almost 800 million years. Carbonate compounds and carbon-rich rocks reacted under the heat and pressure of impact to form carbon dioxide. Ice quickly melted into water. Ammonia, common in the icy tail of comets, reacted to form nitrogen (Figure 17.1). Comets, meteoroids, and asteroids imported Earth's volatiles from outer parts of the Solar System.
The modern atmosphere is mostly
gas, but also contains droplets of liquid water and suspended particles of dust. The gaseous composition of dry air is roughly 78 percent nitrogen, 21 percent oxygen, and 1 percent other gases (Figure 17.5). Nitrogen, the most abundant gas, does not react readily with other substances. Oxygen, though, reacts chemically as fires burn, iron rusts, and plants and animals respire. Carbon dioxide, which by some models formed as much as 80 percent of Earth's early atmosphere, is a trace gas in the modern atmosphere, with a concentration of only 0.035 percent. Composition of the modern atmosphere: 78% Nitrogen, 21% Oxygen and 1% of other gases. The types and quantities of gases, water vapor, droplets, and dust vary with both location and altitude. In a hot, steamy jungle, air may contain 5 percent water vapor by weight, whereas in a desert or cold polar region only a small fraction of a percent may be present. If you sit in a house on a sunny day, you may see a sunbeam passing through a window. The visible beam is light reflected from tiny specks of suspended dust. Clay, salt, pollen, bacteria, viruses, bits of cloth, hair, and skin are all components of dust. People travel to the seaside to enjoy the "salt air." Visitors to the Great Smoky Mountains in Tennessee view the bluish, hazy air formed by sunlight reflecting from pollen and other dust particles. Within the past century, humans have altered the chemical composition of the atmosphere in many different ways. We have increased the carbon dioxide concentration by burning fossil fuels and igniting wildfires. Factories release chemicals into the air—some benign, others poisonous. Smoke and soot change the clarity of the atmosphere. These changes are discussed in the section on "Air Pollution."
Sulfur and nitrogen oxides impair
lung function, aggravating diseases such as asthma and emphysema. They also affect the heart and liver and have been shown to increase vulnerability to viral infections such as influenza. Acid rain corrodes metal and rock. Limestone and marble are especially susceptible because they dissolve rapidly in mild acid. In the United States, the cost of damage and deterioration to buildings and building materials caused by acid precipitation is estimated at several billion dollars per year. Figure 17.15A is a map of average pH for rainfall across the conterminous United States in 1992. The northeastern United States was particularly susceptible to acid rain because it was downwind from the interior "Rust Belt," the large region of the northern interior United States in which large-scale industry had developed since the nineteenth century (Figure 17.15B). More recently, a 2016 study involved the analysis of 27 sites in the northeastern United States and eastern Canada over a period ranging between 8 and 24 years. The results show that long-term increases in pH occurred in the O and B soil horizons at most sites, trends that the study's authors attributed to reversal of the effects of acid rain on North American soils. (A) Annual average measurements of the pH of precipitation during 1992. The highest levels of acidity in precipitation occurred in the northeastern United States, which is generally downwind of the industrial "Rust Belt." (B) Effects of acid rain damage in North Carolina.
As described in section
on "Atmospheric Temperature," solar energy breaks apart oxygen molecules O2 in the stratosphere, releasing free oxygen atoms (O). The free oxygen atoms combine with oxygen molecules to form ozone O3. Ozone absorbs high-energy ultraviolet light. This absorption protects life on Earth because ultraviolet light causes skin cancer, inhibits plant growth, and otherwise harms living tissue. In the 1970s, scientists learned that organic compounds containing chlorine and fluorine, called chlorofluorocarbons (CFCs), and compounds containing bromine and chlorine, called halons, rise into the upper atmosphere, react with, and destroy ozone (Figure 17.18). CFCs have no natural source, but were entirely synthesized for such diverse uses that included cooling agents in almost all refrigerators and air conditioners, propellants in aerosol cans, and as an expanding agent in styrofoam coffee cups and some polystyrene building insulation manufactured before the mid-1980s. CFCs destroy the ozone layer in a three-step reaction. Step 1: CFCs rise into the stratosphere. Ultraviolet radiation breaks the CFC molecules apart, releasing chlorine atoms. Step 2: Chlorine atoms react with ozone, O3, to destroy the ozone molecule and release oxygen, O2. The extra oxygen atom combines with chlorine to produce ClO. Step 3: The ClO sheds its oxygen, producing another free chlorine atom. Thus, chlorine is not used up in the reaction, and one chlorine atom reacts over and over again, to destroy many ozone molecules. In 1985, scientists observed an unusually low ozone concentration in the stratosphere over Antarctica, a phenomenon called the ozone hole. Since it was discovered, the low concentration of ozone over Antarctica declined further until, by 1993, it was 65 percent below normal over 23 million square kilometers, an area almost the size of North America. Research groups also reported significant increases in ultraviolet radiation from the Sun at ground level in the region. In addition, scientists recorded ozone depletion in the Northern Hemisphere. In March 1995, ozone concentration above the United States was 15 to 20 percent lower than during March 1979. Data on global ozone depletion persuaded the industrial nations of the world to limit the use of CFCs and other ozone-destroying compounds. In a series of international agreements signed between 1978 and 1992, many nations of the world agreed to reduce or curtail production of compounds that destroy atmospheric ozone. Most industrialized countries stopped production of CFCs on January 1, 1996. The international bans have had positive results. The concentration of ozone-destroying chemicals peaked in the troposphere (lower atmosphere) in 1994 and has been declining ever since, although CFCs are projected to remain in the atmosphere for over 100 years. The CFCs and halons already in the atmosphere break down slowly, but at least the measured concentrations of these ozone-destroying chemicals are declining. In a review article published in May 2006, the authors concluded that "ozone abundances have at least not decreased...for most of the world," although natural cycles make it difficult to judge the relative effects of human and natural influences. In summary, they write, "it is therefore unlikely that ozone will stabilize at levels observed before 1980, when a decline in ozone concentrations was first observed" (Figure 17.19). As of 2016, eleven different CFC compounds known to be completely synthetic still occur in Earth's atmosphere (Table 17.2). (A) A satellite image of stratospheric ozone concentrations over Antarctica in 1994. (B) Similar data for 2004. In both images, dark purple shows the lowest ozone concentration.
Tropical Storm Toraji
on September 2, 2013. The eye of the storm is in the East China Sea. To the west of the storm's eye, the light blue, shallow continental shelf of eastern mainland China is visible, while suspended sediment muddies the water closer to the shore. Tropical Storm Toraji spawned tornadoes in Japan, the outline of which is shown in the top center of the image. To the east is the deep blue water of the Pacific Ocean. Nearly every multicellular organism needs oxygen to survive. If the oxygen abundance in the atmosphere were to drop below 44 percent of its current value, life on Earth as we know it would perish. If oxygen is essential to most life, would we be better off if we had an even greater supply? The answer is yes, to a limit. Even at sea level, athletes can enhance their performance by breathing a small amount of bottled oxygen. But, paradoxically, too much oxygen is poisonous. If you breathe air that has 55 percent or more oxygen than is found at sea level, your body metabolism becomes so rapid that essential molecules and enzymes decompose. In addition, fires burn more rapidly with increased oxygen concentration. If the oxygen level in the atmosphere were to rise significantly, extremely large wildfires would engulf the planet, altering ecosystems as we know them. Earth is the near-perfect size and distance from a stable and medium-temperature star to permit optimal atmospheric conditions for life. But Earth's atmospheric composition and temperature are not determined solely by planetary size and distance from the Sun. Rather, our planet's environment is finely regulated by interactions among Earth systems. Over the past 4 billion years, the Sun's output has slowly increased (although there were numerous fluctuations during this period). However, Earth's atmospheric temperature has remained remarkably constant because of interactions among Earth's systems. In this chapter, we will explore these interactions and see how they led to the structure and dynamics of Earth's atmosphere today.
Scientists have
revealed a nearly continuous record of rocks of different ages, from 4.04 billion years ago to the present. Therefore, when they study the history of Earth's continental crust, they can analyze the chemical composition of these ancient rocks for clues. Unfortunately, there are no samples of very old atmospheres. So how can we determine atmospheric composition millions to billions of years ago? Many gaps in understanding exist, but the history described next comes from information derived from computer modeling as well as the direct study of rocks. Modeling involves calculations about how atmospheric gases would have behaved under the presumed environment of early Earth. To test these models, scientists study the geochemistry of ancient rocks. Rocks react with water and air, and the nature of these reactions depends on the geochemistry of the entire system. By studying the rocks that existed in a specific time period, scientists deduce the other components of the system that would have produced those reactions. For example, as we will discuss, iron reacts with oxygen to produce iron oxides. Thus, if we find iron oxides in certain types of sedimentary rocks that formed 2.6 billion years ago, we deduce that oxygen must have been present in the air and water at that time.
As explained earlier
sulfur and nitrogen oxides are released when coal and petroleum burn. These oxides are also released when metal ores are refined. In moist air, sulfur dioxide reacts to produce sulfuric acid and nitrogen oxides react to form nitric and nitrous acid. These strong atmospheric acids dissolve in water droplets and fall as acid precipitation, also called acid rain (Figure 17.14). Acid rain develops from the addition of sulfur compounds to the atmosphere by industrial smokestacks. Acidity is expressed on the pH scale. A solution with a pH of 7 is neutral, neither acidic nor basic. On a pH scale, numbers lower than 7 represent acidic solutions and numbers higher than 7 represent basic ones. For example, soapy water is basic and has a pH of about 10, whereas vinegar is an acid with a pH of 2.4. Rain reacts with carbon dioxide in the atmosphere to produce a weak acid. As a result, natural rainfall has a pH of about 5.7. However, in the "bad old days" before the Clean Air Act was properly enforced, rain was much more acidic. A fog in Southern California in 1986 reached a pH of 1.7, which approaches the acidity of toilet bowl cleaners.
Recall from Chapter 15
that a volatile compound is one that evaporates readily and therefore easily escapes into the atmosphere. Whenever chemicals are manufactured or petroleum is refined, some volatile by-products escape into the atmosphere. When metals are extracted from ores, gases such as sulfur dioxide are released. When pesticides are sprayed onto fields and orchards, some of the spray is carried off by the wind. When you paint your house, the volatile parts of the paint evaporate into the air. As a result of all these processes, tens of thousands of volatile compounds are present in polluted air: some are harmless, others are poisonous, and many have not been studied. Consider the case of dioxin. Very little dioxin is intentionally manufactured. It is not an ingredient in any herbicide, pesticide, or other industrial formulation. You cannot buy dioxin at your local hardware store or pharmacy. Dioxin forms as an unwanted by-product in the production of certain chemicals and when specific chemicals are burned. For example, in the United States today, garbage incineration is the most common source of dioxin. When a compound containing chlorine, such as the plastic polyvinyl chloride (PVC), is burned, some of the chlorine reacts with organic compounds to form dioxin. The dioxin then goes up the smokestack of the incinerator, diffuses into the air, and eventually falls to Earth. Cattle eat grass lightly dusted with dioxin and store the dioxin in their fat. Humans ingest the compound mostly in meat and dairy products. The EPA estimates that the average U.S. citizen ingests about 0.0000000001 gram (100 picograms) of dioxin in food every day. Although this is a minuscule amount, the EPA has argued that dioxin is the most toxic chemical known and that even these low background levels may cause adverse effects such as cancer, disruption of regulatory hormones, reproductive and immune system disorders, and birth defects. Others disagree. The Chemical Manufacturers Association has written: "There is no direct evidence to show that any of the effects of dioxins occur in humans in everyday levels." A 2016 study published in the Journal of Epidemiology presented results from an 18 year-long study in which dioxin levels were measured from the milk of first-time Japanese mothers. The study showed that dioxin levels fell between 1998 and 2014 from 20.8 to 7.2 picograms of dioxin toxic equivalence in milk fat, despite a trend toward Japanese women having children later in life during this study period. Because dioxin accumulates in humans through time, milk from older mothers has higher dioxin levels. Thus, the lower overall dioxin levels over the course of the study suggests that concentrations in Japanese first-time mothers fell over the study period. No one knows whether very small doses of potent poisons ingested over long periods of time are harmful. Environmentalists argue that it is "better to be safe than sorry" and that therefore we should reduce ambient concentrations of volatiles such as dioxin. This argument has helped spur the movement toward increased recycling and decreasing incineration of waste plastic as a means of disposal. Others counter that because the harmful effects of compounds like dioxin are unproven, we should not burden our economy with the costs of control.
Ever since
the first cave dwellers huddled around a smoky fire, people have introduced impurities into the air. The total quantity of these impurities is minuscule compared with the great mass of our atmosphere and with the monumental changes that occurred during the evolution of the planet. Yet air pollution remains a significant health, ecological, and climatological problem for modern industrial society. In 1948, Donora was an industrial town with a population of about 14,000 located 50 kilometers south of Pittsburgh, Pennsylvania. One large factory in town manufactured structural steel and wire, and another produced zinc and sulfuric acid. During the last week of October 1948, dense fog settled over the town. But it was no ordinary fog; the moisture contained pollutants from the two factories. After four days, visibility became so poor that people could not see well enough to drive, even at noon with their headlights on. Gradually at first, and then in increasing numbers, residents sought medical attention for nausea, shortness of breath, and constrictions in the throat and chest. Within a week, 20 people had died and about half of the town was seriously ill. Other incidents similar to what happened in Donora occurred worldwide. In response to the growing problem, the United States enacted the Clean Air Act in 1963. As a result of the Clean Air Act and its amendments, total emissions of air pollutants have decreased and air quality across the country has improved (Figure 17.11). It is even more encouraging to note that this decrease in emissions has occurred at a time when population, energy consumption, vehicle miles traveled, and gross domestic product (GDP) have increased dramatically (Figure 17.12). Donora-type incidents in the United States have not been repeated. Smog has decreased, and rain has become less acidic. Yet some people believe that we have not gone far enough and that air pollution regulations should be strengthened further. Measurements of air pollutant concentrations in the United States between 1990 and 2015. This plot, from the U.S. EPA, shows average national concentrations of each air pollutant has fallen through the period, reflecting increased awareness of air quality and implementation of air quality measures. Between 1970 and 2015, a continuous decrease in the aggregate emissions of six common pollutants occurred while gross domestic product increased along with population and total vehicle miles traveled. The six air pollutants included in this plot are: particulate matter of 2.5 and 10 microns diameter, sulfur dioxide, nitrous oxides, volatile organic compounds, carbon monoxide, and lead. Sources and types of air pollution are listed in Figure 17.13 and discussed in the following section. (A) Sources of air pollution in the United States. (B) Types of air pollutants in the United States. (Although carbon dioxide is a greenhouse gas, it is not listed as a pollutant because it is not toxic.)
As explained previously
the first living organisms may have formed by the accretion of complex abiotic (nonliving) organic molecules formed in early Earth's surficial environments. However, complex organic molecules, such as those postulated to have led to the first organisms, become oxidized and destroyed in an oxygen-rich environment. (This oxidation process is analogous to slow burning.) These chemical reactions suggest that if large amounts of oxygen were present in Earth's early atmosphere, the abiotic precursors to living organisms could not have formed. A geochemist can determine whether a rock formed in an oxygen-rich or an oxygen-poor environment by analyzing the mineral assemblage and geochemistry of the rock and applying what is currently known about the thermal and chemical conditions needed to form that suite of minerals. Recent studies of Earth's oldest rocks indicate that the atmospheric oxygen concentration in early Precambrian time was extremely low. These results suggest that the molecules necessary for the emergence of living organisms—namely methane CH4 , ammonia NH3 , hydrogen H2, and water H2O —would have been preserved in the primordial atmosphere. Although life could not have emerged in an oxygen-rich environment, complex multicellular life requires an oxygen-rich atmosphere to survive. How, then, did oxygen become abundant in our atmosphere? The world's earliest organisms probably were chemoautotrophs that obtained their energy from reactions with minerals such as iron and sulfur in an extremely inefficient process. (Recall from Chapter 15 that hemoautotrophic bacteria are common today around vents of extremely hot hydrothermal water emanating from parts of the Mid-Oceanic Ridge.) Later, organisms subsisted, in part, by eating each other. But these food chains were limited because there were only a few organisms on Earth. A crucial step in evolution occurred when primitive bacteria evolved the ability to harness the energy in sunlight and produce organic tissue. This process, known as photosynthesis, is the foundation for virtually all modern life. During photosynthesis, organisms convert carbon dioxide and water to organic sugars. They release oxygen as a by-product. In 1972, an English chemist named James Lovelock hypothesized that the oxygen produced by primitive organisms gradually accumulated, creating the modern atmosphere. By Late Precambrian time, atmospheric oxygen concentration had reached the critical level needed to sustain efficient metabolism. As a result, multicellular organisms evolved and the biosphere as we know it was born. Lovelock was so overwhelmed by the intimate connection between living and nonliving components of Earth's systems that he likened our planet to a living creature, which he called Gaia (Greek for "Earth"). The Lovelock hypothesis, now over 45 years old, remains generally accepted, but scientists are still investigating many of the details. For example, blue-green algae called cyanobacteria began producing oxygen 2.7 billion years ago, but appreciable quantities of oxygen didn't begin to appear in the atmosphere until 2.4 billion years ago, in the First Great Oxidation Event (Figure 17.2), when the concentration of atmospheric oxygen rose abruptly. The oxygen content of the atmosphere slowly and steadily increased after photosynthesizing cyanobacteria began to release oxygen 2.7 billion years ago. The production of oxygen by cyanobacteria beginning 2.7 billion years ago started slowly, and for about 700 million years thereafter, atmospheric oxygen concentration remained far below today's level. At times during this 700-million-year period, Earth's atmosphere may have been anoxic, meaning there was no oxygen. These abrupt variations in the concentration of oxygen in Earth's earliest atmosphere—from periods of anoxia on the one hand to the First Great Oxidation Event 2.4 billion years ago on the other—strongly suggest that the presence of oxygen was driven by a mechanism involving a chemical threshold. To understand the mechanism behind this threshold reaction, we must study atmosphere—geosphere and atmosphere—biosphere systems interactions.
