The energy to heat the Earth comes mainly from the sun. Based on the Earth’s distance from the sun and the amount of solar radiation the sun emits, the average temperature on Earth should be around -18C or 0F. Venus (figure 3.1) is another planet whose temperature is inconsistent with its distance from the Sun. Considered by many to be the brightest and most beautiful in the night sky, Venus has an average temperature on Venus would be 100C. The energy processes that contribute to the Earths energy balance appear in figure 3.2. The Earth receives nearly all the energy from the Sun primarily ultraviolet, visible, and infrared radiation. Some of this incoming radiation is reflected back into space by the dust and aerosol particles that are suspended in our atmosphere (25%). Other parts of this incoming radiation are reflected by the surface of the Earth itself, especially those regions white with snow or sea ice (6%). Thus, 31% of the radiation is  received from the sun is reflected. The remaining 69% of the radiation from the Sun absorbed, either by the atmosphere (23%) or by the land masses and oceans (46%). The 60% of the radiation that is absorbed by the atmosphere, either directly from the sun (23%) or from the Earths surface (37%). 46% of the suns radiation that is absorbed and eventually emitted by the Earth, 37% is absorbed in  the atmosphere prior to its emission into space. If you have ever parked your car outside on a hot day, you realize firsthand how greenhouse gases work. The Greenhouse effect is the natural process by which atmosphere gases trap a major portion (about 80%) of the infrared radiation by the Earth. The atmosphere of Venus acts the same, however it traps more heat. This is because it is made up of nearly 96% carbon dioxide, which we see us a far greater concentration than in the Earth’s atmosphere. Carbon dioxide is a greenhouse gas. Greenhouse gases are those gases capable of absorbing and emitting infrared radiation, thereby warming the atmosphere. In addition to CO, other examples include water vapor, methane, nitrous oxide, ozone, and chlorofluorocarbons. In our energy balance discussion, we showed that 80% of the Earth’s absorbed solar radiation is emitted into the atmosphere. The exchange of energy between the earth atmosphere, and space results in  a steady state and a continuous average temperature of earth. The term enhanced greenhouse effect refers to the process in which atmospheric trap and return more than 80% of the energy radiated by the Earth. The popular term global warming often is used to describe the increase in average global temperatures that results from an enhanced greenhouse effect. Why is the amount of greenhouse gases in the atmosphere increasing? One explanation considers anthropogenic influences on the environment, which stem from human activities such as industry, transportation, mining, and agriculture. These activities require carbon based fuels, which produce carbon dioxide when burned


Methane, nitrous oxide, chlorofluorocarbons, and even ozone all take part in trapping heat in the atmosphere. Our level of concern regarding each of these gases is related to their concentration in the atmosphere but also to other important characteristics. The global atmospheric lifetime characterizes  the time required for a gas added to the atmosphere to be removed. This is quantified by the global warming potential, a number that represents the relative contribution of the molecule of the atmospheric gas to global warming. The GWP of carbon dioxide is assigned the reference value 1; all other greenhouse gases are indexed with respect to it. Table 3.3 lists four greenhouse gases, their main sources, and their important properties in the climate change conversation. The current atmospheric concentration of CH4 is about 50 times lower than that of CO2, but as an infrared absorber, methane is about 20 times more efficient than carbon dioxide. By comparison, carbon dioxide is much less reactive. The primary removal mechanisms for CO2 are dissolution in oceans, photosynthesis by plants, and the much longer process of mineralization into carbonate rocks. Methane emissions arise from both natural and human sources. These marshy habitats are perfectly suited for anaerobic bacteria, those that can function without the use of molecular oxygen. Methane is also released from the oceans, where a substantial amount of it appears to be trapped in “cages” made of water molecules. Australia’s Commonwealth Scientific and Industrial Research Organization has taken a series of ocean core drillings to gather evidence about methane hydrates and their role in global warming (figure 3.23). Termites are another natural source of methane. Not only can they inflict direct damage to homes, but they also add to greenhouse gas concentration. The majority human source of Ch4 is agriculture, with the biggest culprits being rice cultivation and the raising of livestock. Additional agriculture CH4 comes from an increasing number of cattle and sheep. The other main anthropogenic source of methane originates from out extraction of fossil fuels. In the green box it talk about methane concentrations. Another gas that contributes to global warming is nitrous oxide also known as laughing gas. The majority of N2O molecules in the atmosphere comes from the bacterial removal  of nitrate ion from soils, followed by removal of oxygen. In table 3.4 the other differences are summarized. As a result methane contributes a lot to the green house gases. 


Humans are not the only creatures bearing the costs of acidic precipitation. Organisms in the worlds surface waters experience a change in environment when acidic precipitations fills lakes and streams. If the pH is lowered below 6.0, fish and other aquatic life are affect (figure 6.26). Numerous studies have reported the progressive acidification of lakes and rivers in certain geographic region. In southeastern Ontario, the average pH of lakes is now 5.0, well below the pH of 6.5 required for a healthy lake. Many areas of the midwestern US have no problem with acidification of lakes and streams, When acidic precipitation falls or runs off into a lake, the pH of the lake drops unless the acid is neutralized. The capacity of a lake or other body of water to resist a decrease in pH is called it Acid-neutralizing capacity. The lakes and streams also have a relatively high concentration of calcium and hydrogen carbonate ions. This occurs as a result of the reaction of limestone with carbon dioxide and water (equation 6.30). One level of complexity is added by annual variations. A surge of acidity may enter the waterways at just the time when fish are spawning or hatching and are more vulnerable.  U.S SO2 emissions have been declining in recent years, and we have seen a corresponding decrease in the sulfate ion concentrations in the lakes of the Adirondacks. Even though NOx emissions have remained fairly constant, the concentration of nitrates in the Adirondacks is increasing in more lakes than not. It appears that nitrogen saturation have occurred in the surrounding vegetation, with more of the acidity ending up in the lakes. The soil in the region of these lakes most likely has lost some of its acid neutralizing  capacity. 


The causes of haze are well understood, but they differ from region to region. In large cities such as Beijing, China (figure 6.24), coal-burning power plants produce the smoke and participate matter that in turn create the haze. Although both contribute to haze, for the purposes of illustrating acid deposition we focus on the latter. Sulfur dioxide is a colorless, so thus gas is not what we are peering through as “haze”. The first step in this reaction of SO2 with oxygen to form SO3, as we saw earlier  in equation 6.18. Sulfur trioxide is also a colorless gas, but it has the property of being hygroscopic, it readily absorbs water from the atmosphere and retains it. Many molecules of sulfuric acid form a tiny droplet, and many of these droplets of sulfuric acid then coalesce to produce larger droplets. These droplets for an aerosol with droplets about a micrometer u diameter. The aerosols of sulfuric acid, which can persist for several days, can travel hundreds of miles downwind. Recall that sulfuric acid ionizes to produce H+ etc. But these acidic aerosols may reacts with the base to produce salts that contain the sulfate ions. Haze is most pronounced in summer when there is most sunlight to accelerate the photochemical reactions leading to the production of sulfuric acid. IN the US, the Clean Air Act of 1970 and it subsequents amendments included provisions to improve the visibility on our national parks. When you can see haze on the horizon, you are most likely also breathing it. Once inhaled, the acidic droplets attack the sensitive tissue of your lungs. People with asthma, emphysema, and cardiovascular disease are the most sensitive. The low cost of electricity from burning coal fails to take in to account the costs to the community of the health effects resulting from NOx and SO2 emissions. A decrease in acidic aerosol levels would benefit many, both in real dollars and in quality of life. Historically, polluted air has exacted a huge price. One of the worst recorded instances of pollution related respiratory illness occurred in London 1952. Due to unusual weather conditions, a deep layer of fog developed that trapped all the smoke and pollutants for five days. The deadly aerosol caused more than 4000 deaths. Again a layer of fog trapped industrial pollutants close to  ground. By noon, the skies had darkened with a choking  aerosol of fog and smoke (Figure 6.25). During the fog, 17 people died. Although acidic fogs can be immediately hazardous to ones health, public concern is growing over the indirect effects of acid deposition. The savings would come principally from reduced costs to treat pulmonary diseases. 


Although acid rain does not affect all metals, unfortunately iron is one that it does Rods of steel are used to strengthen concrete building and road ways. In table 6.2, you can see the effects of acid rain and benefits of acid rain reduction. The problem with iron is that it rusts, as represented by this chemical equation (6.26). Rusting is a slow process. Iron combines rapidly with oxygen only if you heat or ignite it. Equation 6.26 is the overall equation for the process. The role of H+ is evident in equation 6.27. Even pure water has a sufficient concentration of H+ to promote slow rusting. In the second step, the aqueous Fe2+ further reacts with oxygen. Adding the two steps together give equation 6.26. Because iron is inherently unstable when exposed to the natural environment billions of dollars are spent annually to protect exposed iron and steel in bridges, car, buildings, and ships. Coating iron with a thin layer of a second metal such as chromium or zinc means of protection. Automobile paint can be spotted or pitted by acid deposition. To prevent this damage, automobile manufacturers now use acid resistant paints. Acidic rain also damages statues and monuments made up of marble. Figure 6.22 shows a recognizable, but much deteriorated statue of George Washington. Some limestone tombstones are no longer legible. Many priceless and irreplaceable marble statues and buildings are being attacked by airborne acids (figure 6.23). In the green box, 6.23 it tells you about the deterioration and damage. 


Having identified SO2 and NOx as the two major contributors to acid precipitation, we now examine their production over time and strategies for controlling anthropogenic emissions. Some volcanos steadily emit SO2, and under certain conditions, they produce beautiful crystals of elemental sulfur (Figure 6.20). Marine organisms produce dimethyl sulfide gas as a by product. The dimethyl sulfide enters the troposphere and reacts with OH to form SO2. NO is formed any time that high heat causes nitrogen and oxygen gases to react. Anthropogenic SO2 and NOx emissions outpace natural emissions. The amount of sulfur added to the atmosphere by humans is twice that form volcanoes, oceans, and other natural resources. The amount of nitrogen added as NOx by humans is roughly four times that of natural sources such as lighting and the bacteria found in soils. Consult figures 6.14 and 6.16 to review the sources of these emissions. The levels of these two pollutants have changed dramatically over time. Emissions of both pollutants reached their highest levels in the 1970s. Figure 6.21 shows NOx and So2 emissions in the United States. Globally, the levels of SO2 and NOx also are charging over time. Because the NOx emissions originate from millions of small unregulated, and mobile sources, they are difficult to track. Emissions of SO2 can be estimated with a reasonable degree of accuracy. To get an estimate, researchers start with the amount of fossil fuels, and finally subtract our exports. Metal refining is a bit trickier to estimate, as the amount of sulfur released depends on the technologies used. One such estimate publishes in 2011, a decline in world SO2 emissions over the past two decades. Today, the continent of Asia leads in SO2 emissions. IN 1970, the united states emitted about 30 million tons of sulfur dioxide and china about 7 million tons. To meet the massive demand for electricity, Chinas coal output more than doubled in the first decade of this century. India appears to be headed down the same path as energy needs hamper its economic growth. Such technologies that produce far less SO2 and NOx coupled with conservation, are essential for our planet both now and in the future. 


At first glance it might not be obvious how these thousands of vehicles contribute to acid precipitation. Gasoline burns to form CO2 and H2O. But gasoline contains with small amounts of CO, unburned hydrocarbons. Nitrogen oxides already have been identified as contributors to acid rain, but gasoline does not contain nitrogen. Therefore, logic and chemistry assert that nitrogen oxides cannot be formed from burning gasoline. Recall from Chapter 1 that with sufficient energy, nitrogen and oxygen combine to form Nitrogen monoxide (equation 6.19). The reaction of N2 with O2 is not limited to automobile engines. The same reaction occurs when air is heated in the furnace of a coal fired electric power plant. IN the US, the combustion of fossil fuels in electric power plants accounts for over a third of the nitrogen oxide emission (figure 6.16). In the early 1990s, a green chemistry solution to reducing NO emissions and energy consumption was introduced. Once formed, nitrogen monoxide is highly reactive (equation 6.20). The reactive intermediate species A, A’, and A”, present in trace amounts, are synthesized from the VOC molecules. Nitrogen dioxide is a highly reactive, poisonous, red-brown gas with a nasty odor. A key player is the hydroxyl radical. Once formed in the atmosphere, the hydroxyl radical can rapidly react with nitrogen dioxide to yield nitric acid (equation 6.21). 


When carbon is burned with plenty of oxygen, it forms carbon dioxide and liberates large amounts of energy, which of course is the reason for it burning. You can see the equation in (6.16). Coal also contains varying amounts of sulfur. Plants, like other living things, contain sulfur, so some of the sulfur in coal can be traced to this ancient vegetation. However, most of the sulfur in coal came from the sulfate ion as an oxygen source, removing the oxygen and releasing the sulfate ion. The sulfide ion became incorporated into coal formed in freshwater peat bogs has a lower sulfur content. We can approximate its composition with the chemical formula C135H96O9NS. Burning sulfur in air produces sulfur dioxide, a poisonous gas with an unmistakeable choking odor (figure 6.13). You have see the other part of the equation (6.17). Because the sulfur content of coal varies, burning coal produces sulfur dioxide in varying amounts. When coal is burned, the sulfur dioxide produced goes up the smokestack along with the carbon dioxide, what vapor, and small amounts of metal oxide ash. Once in the atmosphere, SO2 can react with oxygen to form sulfur dioxide (6.18). This reaction, although extremely slow, is accelerated by the presence of tiny particles such as the ash that goes up the stack along with the SO2. Once SO3 is formed, it reacts rapidly with water vapor in the atmosphere to form sulfuric acid (equation 6.12). Although pathways exist for the reaction of SO2 with oxygen to form SO3, the majority of SO2 in the atmosphere directly contributes to acid rain by forming sulfurous acid. A chemical calculation can help us better appreciate the vast quantities of SO2 produced by coal-burning power plants. You can see the equation in the diagrams below. The mass of SO2 is equivalent to 40,000 metric tons. The connection between burning coal and sulfur dioxide emissions in the US is evident in Figure 6.14. Transportation is responsible only for a small percent of the emissions because the gasoline and diesel fuel contain relatively low amounts of sulfur. Industrial processes account for the remainder of the emissions. The smokestacks is tall so that the emissions that still remain are carried away from any sudbury by the prevailing winds (Figure 6.15). In the section following, we relate the bigger stories of nitrogen chemistry. 


Carbon dioxide not only has the opportunity to dissolve in the ocean but can dissolve in rain and fresh water everywhere. When CO2 dissolves, the pH drops slightly due to the formation of carbonic acid. We refer to this type of precipitation as acid rain, this is, with a pH below 5. To measure acidity levels anywhere in the world, we need an analytical tool, the pH meter. When the probe is immersed in a sample, the difference in H+ concentration between the solution and the probe creates a voltage across the membrane. The meter measures this voltage and converts it to a pH value (figure 6.8). Although it is straightforward to measure the pH of a rain sample, certain procedures are necessary to ensure accurate results. For instance, the electrode of the pH meter needs to be carefully calibrated. One way to minimize contamination is to fit a rain collection bucket with a lid and a moisture sensor that opens this lid when it starts to rain. This is how samples are collected at the approximately 250 sites of the National Atmospheric Deposition Program/National Trends Network. One bucket is for dry deposition and the other is covered. A sensor opens this bucket when it rains. Due to budgetary constraints, the test sites cannot go in as many places as researchers might like. The relative advantages of widely dispersing the sites versus putting several near each other in a special ecosystem need to by weighed. Rain samples have been collected routinely in the United States and Canada since about 1970. In figure 6.9 you can see The Bondvile Monitoring Station in central Illinois and in figure B the five active monitoring sites. The photographs in figure 6.10 indicate the magnitude of the operation. The center photo shows a set of rain samples waiting to be analyzed, each assigned an alphanumeric label. A small portion of each sample is saved and stored under refrigeration. From these maps, we observe that all rain is slightly acidic. Remember that rain contains a small amount of dissolved carbon dioxide, and that the carbon dioxide dissolved in rain water produces a weakly acidic solution. In Figure 6.11 it shows that the pH of rain samples is well below normal in the eastern third of the United States. Carbon dioxide is not the only source of H+ in rain. You can see the equation in (6.13, a, b, and c). Then you can see another equation with nitric acid and nitrate ions in (6.14 and 6.15). Rain is only one of several ways that acids can be delivered to Earths surface waters. The term acid deposition refers to both wet and dry forms of delivery of acids from the atmosphere to the surface of the Earth. In figure 6.12 it shows the nitrate and sulfate ion content for wet deposition. Acid deposition also includes the “dry” forms of acids that deposit on land and water. 


Ocean water contains small amounts of three chemical species that play a role in maintaining the ocean pH at approximately 8.2. These three species, carbonate ion, the bicarbonate ion, and carbonic acid. All arise from dissolved carbon dioxide (equation 6.3A, b, c). These species also help maintain your blood at a pH of about 7.4. Human industrial activity has rapidly increased the amount of carbon dioxide released into the atmosphere over the past 200 years. As a result, more carbon dioxide is dissolving into the oceans and forming carbonic acid. In turn, the pH of seawater has dropped by roughly 0.1 pH unit since the early 1800s. A decrease of 0.1 pH unit corresponds to a 26% increase in the amount of H+ in seawater. The lowering of the ocean pH due to increased atmospheric carbon dioxide is called ocean acidification. The H+ produced by the dissociation of carbonic acid reacts with carbonate ion in seawater to form the bicarbonate ion. You can see the equation in (6.10 and 6.11). The interaction of carbonic acid , bicarbonate ion, and carbonate ion are summarized in figure 6.6. As carbon dioxide dissolves in ocean water, it forms carbonic acid. Ocean scientist producer that within the next 40 years, the carbonate ion concentration will reach a low enough level that the shells of sea creatures near the ocean surface will begin to dissolve. To date, only a small number of researchers have focused on the effects of thinning shells in sea creatures. However, negative effects on whole ecosystems have been projected. When changes in ocean pH have occurred over a very long period of time, the ocean has been able to compensate. This happens because large collections of sediment at the bottom of the ocean contain massive amounts of calcuim carbonate. Over time, these sediments dissolve to replenish the carbonate lost to reaction with excess H+. Even if the amount of carbon dioxide in the atmosphere were to immediately level off, the oceans would take thousands of years to return to the pH measured in preindustrial times.