2.9

A major cause of stratospheric ozone depletion was uncovered through the masterful scientific sleuthing of F. The analyzed vast quantities of atmospheric data and studied hundreds of chemical reactions. As the name implies, Chlorofluorocarbons are compounds composed of the element chlorine, fluorine, and carbon. Fluorine and chlorine are members of the same chemical group, the halogens. You can see two examples of CFC in the diagram. CFCs do not occur in nature, we humans synthesized them for a variety of uses. As we saw in the previous section, other contributors to the destruction of ozone, such as the -OH or -NO free radical, are formed in the atmosphere both from natural sources and human activities. Given the desirable nontoxic properties of CFCs, they soon were put to other uses. Halons are close cousins of CFCs. Like them, halons are inert, nontoxic compounds that contain chlorine or fluorine. Halons are used as fire suppressants. For better or worse, the synthesis of CFCs has had a major effect on our lives. Because CFCs are nontoxic, nonflammable, cheap, and widely available, they revolutionized air conditioning. In effect, a major demographic shift occurred because CFC-based technology that transformed the economy and business, potential of entire regions of the globe. It has been estimated that an average CCL2F2 molecule can persist in the atmosphere for 120 years before it meets some fate that decomposes it. In figure 2.14 it shows two sets of data from the latitude at which samples were measure. The major effect is a decrease in ozone and an increase in chlorine monoxide as the South Pole approached it. Because equation 2.10 links ClO-. Cl-, and O3, the conclusion is compelling. Figure 2.14 is sometimes described as the “smoking gun” for stratospheric ozone depletion. Not all of the chlorine implicated in stratospheric destruction comes from CFCs. Other chlorinated carbon compounds come from natural sources. Therefore, any natural chlorine-containing substances are washed out of the atmosphere by rainfall long before they can reach the stratosphere. 

2.8

On average, the total O3 concentration is higher the closer one gets to either pole, with the exception of the seasonal ozone “hole” over the Antarctic. The formulation of ozone in the Chapman cycle is triggered when an O2 molecules absorbs a photon of UV light. Therefore, ozone production increases with the intensity of the tradition striking the stratosphere. The period of highest intensity occurs at the equinox, when the sun is directly overhead. There is a slight increase, however in solar power reaching Earth in early June. Extraordinary images of the Earth, such as the one that opens this chapter, are color coded to show stratospheric ozone concentration. Total ozone levels above earths surface are expressed in Dobson units. Check out the dramatic decline in stratosphere ozone levels observed near the South Pole shown in figure 2.12. In recent years, the minimum has been around 100 Du. The major natural cause of ozone destruction, whatever it takes place around the globe, is a series of reactions involving water vapor and its breakdown products. The great majority of the H2O molecules that evaporate from the oceans and lakes fall back to Earths surface. At this altitude, photons of UV radiation trigger the dissociation of water molecules into hydrogen and hydroxyl free radicals. A free radical is a highly reactive chemical species with one  or more unpaired electrons. Water molecules and their breakdown products are not the only agents responsible for natural ozone destruction. It is part of a natural cycle involving compounds of nitrogen, as we will see in chapter 6. Human activities also can alter steady state concentration of NO. Scientist carried out experiment and calculations to predict the effect of a fleet  of SSTs. They concluded that the risks would outweigh the benefits. Even when the effects of water, nitrogen oxides,  and other naturally occuring compounds are included in stratosphere model, the measured ozone concentration is lower than predicted. The stratospheric ozone concentration at midlatitudes has decreased by more than 8% in some cases. 

2.7

The sunlight that reaches Earths surface contains different types of light, including UV-A, UV-B. When the radiation falls on your skin and is absorbed, it set off a chain of events. First, the energy of the UV photons is deposited in the cell. If the energy is high enough, it may cause some of the chemical bonds in nearby molecules to break. In figure 2.9, you saw a representation of the bond-breaking process. Although bonds can break in many different molecules, those broken in the DNA molecule of a skin cell are of the most concern because the DNA may be mutated in a way that leads to cancer. Skin cancer rates are slowly rising in all countries. Figure 2.11 presents the treads in skin cancer in the US. you can also see that the rate of skin cancer is either constant of very slowly increasing. There are ways to decline the rate of getting skin cancer. For example, wearing sunscreen is one way to reduce the risk. Also wearing protective sunblock is another. Sunblock creams reflect the light. The UV light that hits the surface of our planet affects more than the skin and eyes of human, Plants, animals, and microorganisms are also susceptible to damage. Given the harmful effects of too much UV radiation, you can see why the decreasing stratospheric ozone concentrations observed in the 1980s set off planetary alarm bells. 

2.6

UV-C radiation from the Sun is absorbed in the upper atmosphere before it reaches the ground. Both oxygen and ozone absorb light of these wavelength. About 21% go the atmosphere consists  of oxygen O2. If O2 were the only molecule absorbing UV light from the Sun, Earths surface and the creatures that live on it would still be subjected to damaging radiation in the range of 242-320 nm. The O2 plays an important protective role. Recall that the atoms in O2 molecule are connected with a double bond, but each of the bonds in O3 is somewhere between a single and a double bond in length and in strength. Therefore, the photons of a lower energy are sufficient to separate the atoms on O3. In equation 2.4 and 2.5 together with earlier equation 2.1 are part of a set of chemical reactions in the stratosphere. The process is and example of a steady state, a condition in which a dynamic system is in balance so that there is not net change in concentration of major species involved.  The Chapman cycle as shown in figure 2.1o represents the first set of natural steady state reactions proposed for stratosphere ozone. This natural cycle included chemical reactions for both ozone formation and decomposition. Because of the presence of O2 and O3 in the stratosphere, only certain UV wavelengths reach the surface of the Earth. However, these wavelengths still can cause harm.

2.5

The idea that radiation can be described in terms of wave like characters well established and very useful. In Figure 2.8 the energy of the radiating body were the sum of many energy levels of minute but discrete size. Such an energy distribution is called quantized. Albert Einstein suggested that radiation itself should be viewed as constituted of individual bundles of tested radiation itself should be viewed as constituted of individual bundles of energy called photons. The wave model remains useful, even with the development of the quantum theory. The dual nature of radiant energy seems to defy common sense. The two views are linked in a simple relationship that is one of the most important equations in modern science. You can see this equation in figure 2.3. E represents the energy of a single photon. Both symbols H and C represent constants. The symbol h is called Planks constant and c is the speed of light. Using equation 2.3 one can calculate that the energy associated with a photon of UV radiation is approximately 10 million times larger than energy of a photon emitted by your favorite radio station. You body cannot detect them, but you radio can, The energy associated with each of the radio photons is very low. The sun bombers earth with countless photons. The cells of our retinas are tuned to the wavelengths are absorbed. Compared with animals, green plants captured most of their photons in an even narrower region of the spectrum. Photosynthesis is the process through which green plants and some bacteria capture the energy of sunlight to produce glucose and oxygen from carbon dioxide and water. Photons in the UV region of the spectrum are sufficiently energetic to displace electrons from neutral molecules. The interaction of UV radiation with chemical bonds is shown schematically in figure 2.9. 

2.3

Each hydrogen atom has one electron. If two hydrogen atoms bond then two electrons become common property. Each atom effectively has a share in both electrons. The resulting H2 molecule has a lower energy than the sum of the energy in the two individual H atom. The two electrons that are shared constitute a covalent bond. A Lewis structure is a representation of an atom or molecule that shows its outer electron. We first illustrate the procedure with hydrogen fluoride. You can see the structure of this equation in bullet points 1, 2, and 3. A single covalent bond is formed when two electrons are shared between two atoms. Sometimes the nonbonding electrons are removed from a Lewis structure. This result is called a structural formula. The fact that electrons in many molecules are arranged so that every atom shares in eight electrons is called the octet rule. This generalization is useful for predicting Lewis structures and the formulas of compounds. Throughout the reading it provides multiple examples of Lewis structures. Later on the reader gets introduced to double bonds, and even triple bonds.  A triple bond is a covalent linkage made up of three pairs of shared electrons. Triple bonds are even shorter, stronger, and harder to break up than double bonds. As a result, 2.3 gives examples of Lewis structure and how to solve them in an equation. 

2.2

Both O2 and O3 molecules are composed of oxygen atoms. Every atom has a nucleus, a minuscule and highly dense center of an atom composed of protons and neutrons. Protons are positively charged particles, and neutrons are electrically neutral particles. Outside the nucleus are the electrons that define boundary of the atom. An electron has a mass much smaller than that of a proton, but the opposite side. The number of protons in the nucleus determines the identity of the atom. The term atomic number refers to the number of protons in the nucleus. For example, hydrogen contains one proton, which makes its atomic number one. Electrons are sometime pictured as moving orbits about the nucleus. The periodic table lists elements in order of increasing atomic number. The table also has elements arranged so that those with similar chemical properties fall into the same column. The electrons in the innermost level are most strongly attracted by the positively charged protons in the nucleus. The greater distance between an electron and the nucleus, the weaker attraction between them. Each engr. love has a maximum number of electrons that can be accommodated and is particularly stable when fully opposed. The innermost level can hold only two electrons. The second level has a maximum capacity of eight. Out er electrons are found in the highest energy level and help to account for many of the observed trends in the chemical properties. In table 2.2 it shows the 18 elements. the periodic table is a useful guide to electron arrangement. In addition to electrons and protons, atoms also contain neutrons. The one exception is an atom of the most common form of hydrogen, which consists of only one proton in its nucleus. Hydrogen, deuterium, and tritium are examples of isotopes. Isotopes are two or more forms of the same element whose atoms are differ in number of neutrons, and hence in mass. An isotope is identified by its mass number, the sum of the number of protons and neutrons in the nucleus of an atom. All elements have more than one isotope. Each elements atomic mass takes a relative natural abundance of isotopes.

2.1

The ozone and oxygen molecules differ by only one atom. This difference is molecular structure translates to significant differences in chemical properties. One difference is that ozone is far more chemically reactive than O2. Ozone forms both naturally and as a result of human activities. Ozone can be formed from oxygen, but the process requires energy. You can see the equation is figure 2.1. Ozone is reasonably rare in the troposphere, the region of the atmosphere closest to the Earth’s surface. 20 and 100 ozone molecules typically occur for each billion molecules and atoms that make up the air. The stratosphere is where ozone does most of its filtering of some types of ultraviolet light from the sun.Most of our ozone on this plant which is about 90%, is found in the stratosphere. The term ozone layer refers to a designated region in the stratosphere of maximum ozone concentration. Im figure 2.2 its shows the ozone concentration in the troposphere and stratosphere. Scientist continue to measure and evaluate ozone levels using ground observations, weather ballons, and high-flyig aircraft. This process by which ozone protects us from damaging solar radiation involves the interaction of matter and energy from the sun. 

1.3

Indoor air may contain up to a thousand substances at low levels. Indoor air contains some familiar culprits. For example, VOCs, NO, NO2, SO2, Co, ozone, and PM. These pollutants are presented because they came in with the outside air or because they were generated right outside your dwelling. If a pollutant is highly reactive, it does not persist long enough to be transported indoors. For highly reactive molecules such as O3, NO2, and SO2, you expect lower levels indoors. Indoors air is typically 10-30% lower in ozone concentration than outdoor air. Similarly, sulfur dioxide and nitrogen dioxide levels are lower indoors, even though the decrease is not as dramatic as that for ozone. On the other hand, carbon monoxide is a different story. CO has a long enough atmosphere lifetime to move freely in and out through building. Some pollutants are trapped by the filters in the heating or cooling system of the building. People who suffer from seasonal allergies can fin relief indoors. However, gas molecules such as O3, CO, NO2, and SO2 are not trapped by the filters used in most ventilation systems. A building that is air-tight with a limited intake fresh air may have unhealthy levels of indoor air pollutants. Poor ventilation can cause indoor pollutants to reach hazardous levels. Even with good ventilation, indoor activities can compromise air equality. Like tobacco smoke is a serious indoor air pollutant that contains over a thousand chemical substances. Also when you light a candle it depletes the oxygen in the room. Burning candles or incense can generate fumes more rapidly. In figure 1.20, paints and varnishes are a source of VOCs. All oil based paints and varnishes rank higher than water-based ones. Some VOCs are additives that evaporates the paint dries. Coalescent’s are another addictive used in latex paints. Coalescent’s are chemicals added to soften the latex particles. Stricter government regulations on VOC emissions have prompted paint manufacturers to defuse new formulations for latex paints. Advantages of these new coalescent’s is that they are produced from vegetable oils. There is no loss in quality, as the paints formulated with vegetable oil base coalescent’s meet to exceed the performance of traditional paints. In the end, indoors and outdoors there will be pollutants.

1.12

The ozone is a bad actor in the troposphere. The ozone also damages tress and reduces lung function in healthy people. In the absences of sunlight, ozone does not persist for long. After sundown, the ozone levels drop. That statement is from the green box and is depicted in figure 1.18. Unlike the pollutants described in section 1.11, ozone is a secondary pollutant. It is produced from chemical reactions involving one or more other pollutants. . In figure 1.18, the maps show the ozone pollution on a summer day in July 2006. The nitrogen dioxide meets several fates in the atmosphere. The energy provided by sunlight splits one of the bonds into the NO2 molecule. You can see that in the equation (1.13). Then the oxygen atoms produced then can react with oxygen molecules to produce ozone (1.14). Sunlights splits NO2 to release O atoms. These in turn reacts with O2 to form O3. All three are found in nature, but O2 is the least reactive and by far the most abundant. This is because about one fifth of the air we breathe. Oxygen atoms also exist in our upper atmosphere and are even more reactive than ozone. Another example of the tragedy of the commons is Canada’s polluted air. The tragedy arises when a resources is common to all and used by many. Air pollution is a serious international issue. Ozone attacks rubber and so it damages the tires of vehicles that led to its production in the first place.