The Ozone Layer

In this chapter, the following introductory chemistry topics are used:

■ Moles; concentration units including mole fraction

■ Ideal gas law; partial pressures

■ Thermochemistry: AH, AHf} Hess' law

■ Kinetics: rate laws; reaction mechanisms, activation energy, catalysis

Introduction Ozone Layer


The ozone layer is a region of the atmosphere that is called "Earth's natural sunscreen" because it filters out harmful ultraviolet (UV) rays from sunlight before they can reach the surface of our planet and cause damage to humans and other life forms. Any substantial reduction in the amount of this ozone would threaten life as we know it. Consequently, the appearance in the mid-1980s of a large "hole" in the ozone layer over Antarctica represented a major environmental crisis. Although steps have been taken to prevent its expansion, the hole will A young girl applies sunscreen to protect her skin against UV rays from continue to appear each Spring over the Sun. [Source: Lowell George/CORBIS.] the South Pole; indeed, one of the largest, deepest holes in history occurred in 2006. Thus it is important that we understand the natural chemistry of the ozone layer, the subject of this chapter. The specific processes at work in the ozone hole and the history of the evolution of the hole are elaborated upon in Chapter 2. We begin by considering how the concentrations of atmospheric gases are reported and the region of the atmosphere where the ozone is concentrated.

Regions of the Atmosphere

The main components (ignoring the normally ever-present but variable water vapor) of an unpolluted version of the Earth's atmosphere are diatomic nitrogen, N2 (about 78% of the molecules); diatomic oxygen, 02 (about 21%); argon, Ar (about 1%); and carbon dioxide, C02 (presently about 0.04%). (The names of chemicals important to a chapter are printed in bold, along with their formulas, when they are introduced. The names of chemicals less important in the present context are printed in italics.) This mixture of chemicals seems unreactive in the lower atmosphere, even at temperatures or sunlight intensities well beyond those naturally encountered at the Earth's surface.

The lack of noticeable reactivity in the atmosphere is deceptive. In fact, many environmentally important chemical processes occur in air, whether clean or polluted. In the next two chapters, these reactions will be explored in detail when we discuss reactions that occur in the troposphere, the region of the atmosphere that extends from ground level to about 15 kilometers altitude and contains 85% of the atmosphere's mass. In this chapter we will consider processes in the stratosphere, the portion of the atmosphere from approximately 15 to 50 kilometers (i.e., 9-30 miles) altitude that lies just above the troposphere. The chemical reactions to be considered are vitally important to the continuing health of the ozone layer, which is found in the bottom half of the stratosphere. The ozone concentrations and the average temperatures at altitudes up to 50 kilometers in the Earth's atmosphere are shown in Figure 1-1.

The stratosphere is defined as the region that lies between the altitudes where the temperature trends display reversals: The bottom of the stratosphere occurs where the temperature first stops decreasing with height and begins to increase, and the top of the stratosphere is the altitude where the temperature stops increasing with height and begins to decrease. The exact altitude at which the troposphere ends and the stratosphere begins varies with season and with latitude.

Environmental Concentration Units for Atmospheric Gases

Two types of concentration scales are commonly used for gases present in air. For absolute concentrations, the most common scale is the number of molecules per cubic centimeter of air. The variation in the concentration of ozone on the molecules per cubic centimeter scale with altitude is illustrated in Figure 1-la. Absolute concentrations are also sometimes expressed in

FIGURE 1-1 Variation with altitude of (a) ozone concentration (for mid-latitude regions) and (b) air temperature for various regions of the lower atmosphere.

terms of the partial pressure of the gas, which is stated in units of atmospheres or kilopascals or bars. According to the ideal gas law (PV = nRT), partial pressure is directly proportional to the molar concentration n/V, and hence to the molecular concentration per unit volume, when different gases or components of a mixture are compared at the same Kelvin temperature T.

Relative concentrations are usually based on the chemists' familiar mole fraction scale (called mixing ratios by physicists), which is also the molecule fraction scale. Because the concentrations for many constituents are so small, atmospheric and environmental scientists often re-express the mole or molecule fraction as a parts per_value. Thus a concentration of 100 molecules of a gas such as carbon dioxide dispersed in one million (106) molecules of air would be expressed as 100 parts per million, i.e., 100 ppm, rather than as a molecule or mole fraction of 0.0001. Similarly, ppb and ppt stand for parts per billion (one in 109) and parts per trillion (one in 10'2).

It is important to emphasize that for gases, these relative concentration units express the number of molecules of a pollutant (i.e., the "solute" in chemists' language) that are present in one million or billion or trillion molecules of air. Since, according to the ideal gas law, the volume of a gas is proportional to the number of molecules it contains, the "parts per" scales also represent the volume a pollutant gas would occupy, compared to that of the stated volume of air, if the pollutant were to be isolated and compressed until its pressure equaled that of the air. In order to emphasize that the concentration scale is based upon molecules or volumes rather than upon mass, a v (for volume) is sometimes shown as part of the unit, e.g., 100 ppmv or 100 ppmv.

Ozone Variance Altitude

Ozone concentration in units of 1012 molecules/cm3

Ozone concentration in units of 1012 molecules/cm3

FIGURE 1-1 Variation with altitude of (a) ozone concentration (for mid-latitude regions) and (b) air temperature for various regions of the lower atmosphere.

Temperature (°C)

Temperature (°C)

The Physics and Chemistry of the Ozone Layer

To understand the importance of atmospheric ozone, we must consider the various types of light energy that emanate from the Sun and consider how UV light in particular is selectively filtered from sunlight by gases in the air. This leads us to consider the effects on human health of UV light, and quantitatively how energy from light can break apart molecules. With that background, we can then investigate the natural processes by which ozone is formed and destroyed in air.

Absorption of Light by Molecules

The chemistry of ozone depletion, and of many other processes in the stratosphere, is driven by energy associated with light from the Sun. For this reason, we begin by investigating the relationship between light absorption by molecules and the resulting activation, or energizing, of the molecules that enable them to react chemically.

An object that we perceive as black in color absorbs light at all wavelengths of the visible spectrum, which runs from about 400 nm (violet light) to about 750 nm (red light); note that one nanometer (nm) equals Iq-9 meter_ Sybstances differ enormously in their propensity to absorb light of a given wavelength because of differences in the energy levels of their electrons. Diatomic molecular oxygen, 02, does not absorb visible light very readily, but it does absorb some types of ultraviolet (UV) light, which is that having wavelengths between about 50 and 400 nm. The most environmentally relevant portion of the electromagnetic spectrum is illustrated in Figure 1-2. Notice that the UV region begins at the violet edge of the visible region, hence the name ultraviolet. The division of the UV region into components will be discussed later in this chapter. At the other end of the spectrum, beyond the red portion of the visible region, lies infrared light, which will become important to us when we discuss the greenhouse effect in Chapter 6.

An absorption spectrum such as that illustrated in Figure 1-3 is a graphical representation that shows the relative fraction of light that is absorbed by a given type of molecule as a function of wavelength. Here, the efficient light-absorbing behavior of 02 molecules for the UV region between 70 and 250 nm is shown; some minuscule amount of absorption continues beyond 250 nm, but in an ever-decreasing fashion (not shown). Notice that the fraction of light absorbed by 02 (given on a logarithmic scale in Figure 1-3) varies quite dramatically with wavelength. This sort of selective absorption behavior is observed for all atoms and molecules, although the specific regions of strong absorption and of zero absorption vary widely, depending upon the structure of the species and the energy levels of their electrons.




200 280 320 400

Violet Red


100,000 (100 fjm)

Thermal 1R

FIGURE 1-2 The electromagnetic spectrum. The ranges of greatest environmental interest in this book are shown.

Atmospheric Selective Absorbers

FIGURE 1-3 Absorption spectrum of 02. ISource: T. E. Graede! and P.). Crutzen, Atmospheric Change: An Earth System Perspective (New York: W. H. Freeman, 1993).]

FIGURE 1-3 Absorption spectrum of 02. ISource: T. E. Graede! and P.). Crutzen, Atmospheric Change: An Earth System Perspective (New York: W. H. Freeman, 1993).]

Filtering of Sunlight s UV Component by Atmospheric 02 and 03

As a result of these absorption characteristics, the 02 gas that lies above the stratosphere filters from sunlight most of the UV light from 120 to 220 nm; the remainder of the light in this range is filtered by the 02 in the stratosphere. Ultraviolet light that has wavelengths shorter than 120 nm is filtered in and above the stratosphere by 02 and other constituents of air such as N2. Thus no UV light having wavelengths -1 shorter than 220 nm reaches the Earth's surface. This screening protects our skin and eyes, and in fact protects all biological life, from extensive damage by this part of the Sun s output.

Diatomic oxygen also filters some, but not all, of sunlight's UV in the 220-240-nm range. Instead, ultraviolet light in the whole 220-320-nm range is filtered from sunlight mainly by ozone molecules, 03, that are spread through the middle and lower stratosphere. The absorption spectrum of ozone in this wavelength region is shown in Figure 1-4. Since its molecular constitution, and thus its set of energy levels, is different from that of diatomic oxygen, its light absorption characteristics also are quite different.

Ozone, aided to some extent by 02 at the shorter wavelengths, filters out all of the Sun's ultraviolet light in the 220-290-nm range, which overlaps the 200-280-nm region known as UV-C (see Figure 1-2). However, ozone can absorb only a fraction of the Sun's UV light in the 290-320-nm range, since, as you can infer from Figure l-4b, its inherent ability to absorb light of such wavelengths is quite limited. The remaining amount of the sunlight of such wavelengths, 10-30% depending upon latitude, penetrates the atmosphere to the Earth's surface. Thus ozone is not completely effective in shielding us from light in the UV-B region, defined as that which lies from 280 to 320 nm (although different authors vary slightly on the limits for this parameter). Since the absorption by ozone

200 210 220 230 240 250 260 270 280 290 ] 300 Wavejength (nm)

200 210 220 230 240 250 260 270 280 290 ] 300 Wavejength (nm)

Ozone Layer Ks1

305 310 315 Wavelength (nm)

FIGURE 1-4 Absorption spectrum of 03: (a) from 200 to 300 nm and (b) from 295 to 325 nm. Note that different scales are used for the extent of absorption in the two cases. [Sources: (a) Redrawn from M. J. McEwan and L. F. Phillips, Chemistry of the Atmosphere (London: Edward Arnold, 1975). (b) Redrawn from J. B. Kerr and C. T. McElroy, Science 262: 1032-1034. Copyright 1993 by the AAAS.l

305 310 315 Wavelength (nm)

FIGURE 1-4 Absorption spectrum of 03: (a) from 200 to 300 nm and (b) from 295 to 325 nm. Note that different scales are used for the extent of absorption in the two cases. [Sources: (a) Redrawn from M. J. McEwan and L. F. Phillips, Chemistry of the Atmosphere (London: Edward Arnold, 1975). (b) Redrawn from J. B. Kerr and C. T. McElroy, Science 262: 1032-1034. Copyright 1993 by the AAAS.l falls off in an almost exponential manner with wavelength in this region (see Figure l-4b), the fraction of solar UV-B that reaches the troposphere increases with increasing wavelength.

Because neither ozone nor any other constituent of the clean atmosphere absorbs significantly in the UV-A range, i.e., 320400 nm, most of this, the least biologically harmful type of ultraviolet light, does penetrate to reach the Earth's surface. (Nitrogen dioxide gas does absorb UV-A light but is present in such small concentration in clean air that its net absorption of sunlight is quite small.)

The net effect of diatomic oxygen and ozone in screening the troposphere from the UV component of sunlight is illustrated in Figure 1-5. The curve at the left corresponds to the intensity of light received outside the Earth's atmosphere, whereas the curve at the right corresponds to the light that is transmitted to the troposphere (and thus to the surface). The vertical separation at each wavelength between the curves corresponds to the amount of sunlight that is absorbed in the stratosphere and outer regions of the atmosphere.

At top of / /


- J


1 i ll

surface i 1 i 1

300 400

Wavelength (nm)

300 400

Wavelength (nm)

UV light

Visible light

Biological Consequences of Ozone Depletion

A reduction in stratospheric ozone concentration allows more UV-B light to penetrate to the Earth's surface. A 1% decrease in overhead ozone is predicted to result in a 2% increase in UV-B intensity at ground level. This increase in UV-B is the principal environmental concern about ozone depletion, since it leads to detrimental consequences to many life forms, including humans. Exposure to UV-B causes human skin to sunburn and suntan; overexposure can lead to skin cancer, the most prevalent form of cancer. Increasing amounts of UV-B may also adversely affect the human immune system and the growth of some plants and animals.

Most biological effects of sunlight arise because UV-B can be absorbed by DNA molecules, which then may undergo damaging reactions. By comparing the variation in wavelength of UV-B light of differing intensity arriving at

FIGURE 1-5 The intensity of sunlight in the UV and in part of the visible regions measured outside the atmosphere and in the troposphere. [Source: W. L. Chameides and D. D. Davis, Chemical and Engineering News (4 October 1982): 38-52. Copyright 1982 by the American Chemical Society. Reprinted with permission.)

the Earth's surface with the absorption characteristics of DNA as shown in Figure 1-6, it can be concluded that the major detrimental effects of sunlight absorption will occur at about 300 nm. Indeed, in light-skinned people, the skin shows maximum UV absorption from sunlight at about 300 nm.

Most skin cancers in humans are due to overexposure to UV-B in sunlight, so any decrease in ozone is expected eventually to yield an increase in the incidence of this disease. Fortunately, the great majority of skin cancer cases are not the often-fatal (25% mortality rate) malignant melanoma, but rather one of the slowly spreading types that can be treated and that collectively affect about one in four Americans at some point in their lives. The plot in Figure 1-7, which is based on health data from seven countries at different latitudes that therefore receive different amounts of ground-level UV, shows that the rise in the incidence of nonmelanoma skin cancer with exposure to UV is exponential; the reason is that the logarithm of the incidence is linearly related to the UV intensity. For example, the skin cancer rate in Europe is only about half that in the United States.

The incidence of the malignant melanoma form of skin cancer, which affects about 1 in 100 Americans, is thought to be related to short periods of very high UV exposure, particularly early in life. Especially susceptible are fair-skinned, fair-haired, freckled people who burn easily and who have moles with irregular shapes or colors. The incidence of malignant melanoma is also related to latitude. White males living in sunny climates such as Florida or Texas are twice as likely to die from this disease as those in the more northerly states, although part of this greater incidence is probably due to different patterns of personal behavior, such as choice of clothing, as well as to increased UV-B content in the sunlight. Curiously, indoor workers—who have intermittent exposure to the Sun—are more susceptible than are tanned, outdoor workers! The lag period between first exposure and melanoma is 15-25 years. If malignant melanoma is not treated early, it can spread via the bloodstream to body organs such as the brain and the liver.

The phrase full spectrum is sometimes used to denote sunscreens that block UV-A as well as UV-B light. The use of sunscreens that block UV-B, but not UV-A, may actually lead to an increase in melanoma skin cancer, since sunscreen usage allows people to expose their skin to sunlight for prolonged periods without burning. The substances used in sunscreen lotions (e.g., particles of inorganic compounds such as £inc oxide or titanium oxide)

Waves Mitral Regurgitation

FIGURE 1-6 The absorption spectrum for DNA and the intensity of sunlight at ground level versus wavelength. The degree of absorption of light energy by DNA reflects its biological sensitivity to a given wavelength. [Source: Adapted from R. B. Setlow, Proceedings of the National Academy of Science USA 71 (1974): 3363-3366.]

FIGURE 1-6 The absorption spectrum for DNA and the intensity of sunlight at ground level versus wavelength. The degree of absorption of light energy by DNA reflects its biological sensitivity to a given wavelength. [Source: Adapted from R. B. Setlow, Proceedings of the National Academy of Science USA 71 (1974): 3363-3366.]

either reflect or scatter sunlight or absorb its UV component (e.g., water-insoluble organic compounds such as octinoxate— octyl methoxycinnamate—for UV-B absorption and oxyben-zone for UV-A) before it can reach and be absorbed by the skin. Sunscreens were one of the first consumer products to use nanoparticles, i.e., tiny particles only a few dozen or a few hundred nanometers (1CT9 m) in size. Since such particles are so tiny and do not absorb or reflect visible light, the sunscreens appear transparent.

Potential sunscreen compounds are eliminated if they undergo an irreversible chemical reaction when they absorb sunlight, because this would quickly reduce the effectiveness of the application and because the reaction products could be toxic to the skin. Also, the commonly used sunscreen component PABA (p-aminobenzoic acid) is no longer generally used because of evidence that it can itself cause cancer.

The SPF (Sun Protection Factor) of a sunscreen measures the multiplying factor by which a person can stay exposed to the Sun without burning. Thus an SPF of 15 means that he or she can stay in the Sun fifteen times longer than without the sunscreen. To receive that protection, however, the sunscreen must be reapplied at least every few hours.

Because of the long time lag (30-40 years) between exposure to UV and the subsequent manifestation of nonmalignant skin cancers, it is unlikely that effects from ozone depletion are observable as yet. The rise in skin cancer that has occurred in many areas of the world—and that is still occurring, especially among young adults—is probably due instead to greater amounts of time spent by people outdoors in the Sun over the past few decades. For example, the incidence of skin cancer among residents of Queensland, Australia, most of whom are light-skinned, rose to about 75% of the population as lifestyle changes increased their exposure to sunlight years before ozone depletion began. As a consequence of its experience with skin cancer, Australia has led the world in public health awareness of the need for protection from ultraviolet exposure.

In addition to skin cancer, UV exposure has been linked to several other human conditions. The front of the eye is the one part of the human anatomy where ultraviolet light can penetrate the human body. However, the cornea and lens filter out about 99% of UV from light before it reaches the retina. Over time, the UV-B absorbed by the cornea and lens produces highly reactive molecules called free radicals that attack the structural molecules and can produce cataracts. Indeed, there is some evidence that increased UV-B levels give rise to an increased incidence of eye cataracts, particularly among the nonelderly (see Figure 1-8). UV exposure has also been linked to an increase in the rate of macular degeneration, the gradual death of cells in the

100 200 300 400 Intensity of annual ultraviolet radiation (W s/em2)

FIGURE 1-7 Incidence (logarithmic scale) for nonmelanoma skin cancer per 100,000 males versus annual UV light intensity, using data from various countries. [Source: Redrawn from D, Gordon and H. Silverstone, in R. Andrade et al., Cancer of the Skin (Philadelphia: W, B. Saunders, 1976), pp. 405-434.]

FIGURE 1-8 (a) A normal human eye and (b) a human eye with cataract. [Sources: (a) Martin Dohrn/Photo Researchers; (b) Sue Ford/Photo Researchers.]

FIGURE 1-8 (a) A normal human eye and (b) a human eye with cataract. [Sources: (a) Martin Dohrn/Photo Researchers; (b) Sue Ford/Photo Researchers.]

central part of the retina. Increased UV-B exposure also leads to a suppression of the human immune system, probably with a resulting increase in the incidence of infectious diseases, although this has not yet been extensively researched.

However, sunlight does have some positive effects on human health. Vitamin D, which is synthesized from precursor chemicals by the absorption of UV by the skin, is an anticancer agent. Recent research has established that sunlight during the winter is too weak a source of vitamin D synthesis for people living in mid- to high latitudes and that supplementary sources of the vitamin may be advisable. Insufficient vitamin D can reduce the rate of bone regeneration—since the vitamin is required for calcium utilization by the body—and thereby lead to increased fragility among middle-aged and elderly adults. Some controversial research indicates that moderate exposure to the Sun can reduce the incidence of multiple sclerosis and several types of cancer.

Humans are not the only organisms affected by ultraviolet light. It is speculated that increases in UV-B exposure can interfere with the efficiency of photosynthesis, and plants may respond by producing less leaf, seed, and fruit. All organisms that live in the first five meters or so below the surface in bodies of clear water would also experience increased UV-B exposure arising from ozone depletion and may be at risk. It is feared that production of the microscopic plants called phytoplankton near the surface of seawater may be at significant risk from increased UV-B; this would affect the marine food chain for which phytoplankton forms the base. Experiments indicate that there is a complex interrelationship between plant production and UV-B intensity, since the latter also affects the survival of insects that feed off the plants.

Variation in Light s Energy with Wavelength

As Albert Einstein realized, light can not only be considered a wave phenomenon but also to have particle-like properties in that it is absorbed (or emitted) by matter only in finite packets, now called photons. The quantity of energy, E, associated with each photon is related to the frequency, v, and the wavelength, A, of the light by the formulas

Here h is Planck's constant (6,626218 X 1CT34 J s) and c is the speed of light (2.997925 X 108 m s"1). From the equation, it follows that the shorter the wavelength of the light, the greater the energy it transfers to matter when absorbed. Ultraviolet light is high in energy content, visible light is of intermediate energy, and infrared light is low in energy. Furthermore, UV-C is higher in energy than UV-B, which in turn is more energetic than is UV-A.

For convenience, the product he in the equation above can be evaluated on a molar basis to yield a simple formula relating the energy absorbed by 1 mole of matter when each molecule in it absorbs one photon of a particular wavelength of light. If the wavelength is expressed in nanometers, the value of he is 119,627 kj mol 1 nm, so the equation becomes

The photon energies for light in the UV and visible regions are of the same order of magnitude as the enthalpy (heat) changes, AH°, of chemical reactions, including those in which atoms dissociate from molecules. For example, it is known that the dissociation of molecular oxygen into its monatomic form requires an enthalpy change of 498.4 kj mol" :

02-> 2 O AH° = 498.4 kJmoP1

In general, we can calculate enthalpy changes for any reaction by recalling from introductory chemistry that for any reaction, AH0 equals the sum of the enthalpies of formation, AH,0, of the products minus those of the teactants:

In the case of the reaction above,

AH0 = 2 AH? (O, g) - AH° (02, g)

From data tables, we find that AH° (O, g) = +249.2 kj/mol, and we know that AH? (02, g) = 0 since 02 gas is the stablest form of the element. By substitution,

To a good approximation, for a dissociation reaction, AH" is equal to the energy required to drive the reaction. Since all the energy has to be supplied by one photon per molecule (see below), the corresponding wavelength for the light is

A - 119,627 kj moP1 nm/498.4 kj moP1 = 240 nm

Thus any 02 molecule that absorbs a photon from light of wavelength 240 nm or shorter has sufficient excess energy to dissociate.

If energy in the form of light initiates a reaction, it is called a photochemical reaction. The oxygen molecule in the above reaction is variously said to be photochemically dissociated or photochemically decomposed or to have undergone photolysis.

Atoms and molecules that absorb light (in the ultraviolet or visible region) immediately undergo a change in the organization of their electrons. They are said to exist temporarily in an electronically excited state; to denote this, their formulas are followed by a superscript asterisk (*). However, atoms and molecules generally do not remain in the excited state, and therefore do not retain the excess energy provided by the photon, for very long. Within a tiny fraction of a second, they must either use the energy to react photochemically or return to their ground state—the lowest energy (most stable) arrangement of the electrons. They quickly return to the ground state either by emitting a photon themselves or by converting the excess energy into heat that becomes shared among several neighboring free atoms or molecules as a result of collisions (i.e., molecules must "use it or lose it").

x reaction

Consequently, molecules normally cannot accumulate energy from several photons until they receive sufficient energy to react; all the excess energy required to drive a reaction usually must come from a single photon. Therefore, light of 240 nm or less in wavelength can result in the dissociation of 02 molecules, but light of longer wavelength does not contain enough energy to promote the reaction at all, even though certain wavelengths of such light can be absorbed by the molecule (see Figure 1-3). In the case of an 02 molecule, the energy from a photon of wavelength greater than 240 nm can, if absorbed temporarily, raise the molecules to an excited state, but the energy is rapidly converted to an increase in the energy of its own motion and that of the molecules that surround it.

02 + photon (A > 240 nm)-> 02*-> O, + heat

02 + photon (A < 240 nm)-* 02*-» 2 O or 02 + heat


What is the energy, in kilojoules per mole, associated with photons having the following wavelengths? What is the significance of each of these wavelengths? [Hint: See Figure 1-2.J

(a) 280 nm (b) 400 nm (c) 750 nm (d) 4000 nm PROBLEM 1-2

The AH° for the decomposition of ozone into 02 and atomic oxygen is + 105 kj mol-1:

What is the longest wavelength of light that could dissociate ozone in this manner? By reference to Figure 1-2, decide the region of sunlight (UV, visible, or infrared) in which this wavelength falls.


Using the enthalpy of formation information given below, calculate the maximum wavelength that can dissociate N02 to NO and atomic oxygen. Recalculate the wavelength if the reaction is to result in the complete dissociation into free atoms (i.e., N + 2 O). Is light of these wavelengths available in sunlight?

AH? values (kj moF1): N02: +33.2; NO: +90.2; N: +472.7; O: +249.2

Of course, in order that a sufficiently energetic photon supply the energy to drive a reaction, it must be absorbed by the molecule. As you can infer from the examples of the absorption spectra of 02 and 03 (Figures 1-3 and 1-4), there are many wavelength regions in which molecules simply do not absorb significant amounts of light. Thus, for example, because ozone molecules do not absorb visible light near 400 nm, shining light of this wavelength on them does not cause them to decompose, even though 400-nm photons carry sufficient energy to dissociate them to atomic and molecular oxygen (see Problem 1-2). Furthermore, as discussed above, the fact that molecules of a substance absorb photons of a certain wavelength and such photons are sufficiently energetic to drive a reaction does not mean that the reaction necessarily will occur; the photon energy can be diverted by a molecule into other processes undergone by the excited state. Thus the availability of light with sufficient photon energy is a necessary, but not a sufficient, condition for reaction to occur with any given molecule.

Creation of Ozone in the Stratosphere

In this section, the formation of ozone in the stratosphere and its destruction by noncatalytic processes are analyzed. As we shall see, the formation reaction generates sufficient heat to determine the temperature in this region of the atmosphere. Above the stratosphere, the air is very thin and the concentration of molecules is so low that most oxygen exists in atomic form, having been dissociated from 02 molecules by UV-C photons from sunlight. The eventual collision of oxygen atoms with each other leads to the re-formation of 02 molecules, which subsequently dissociate photochemically again when more sunlight is absorbed.

In the stratosphere itself, the intensity of the UV-C light is much lower because much of it is filtered by the diatomic oxygen that lies above. In addition, since the air is denser than it is higher up, the molecular oxygen concentration is much higher in the stratosphere. For this combination of reasons, most stratospheric oxygen exists as 02 rather than as atomic oxygen. Because the concentration of 02 molecules is relatively large and the concentration of atomic oxygen is so small, the most likely fate of the stratospheric oxygen atoms that are created by the photochemical decomposition of 02 is not their mutual collision to re-form 02 molecules. Rather, they are more likely at such altitudes to collide and react with undissociated, intact diatomic oxygen molecules, an event that results in the production of ozone:

Indeed, this reaction is the source of all the ozone in the stratosphere. During daylight hours, ozone is constantly being formed by this process, the rate of which depends upon the amount of UV light and consequently the concentration of oxygen atoms and molecules at a given altitude.

At the bottom of the stratosphere, the abundance of 02 is much greater than that at the top because air density increases progressively as one approaches the surface. However, relatively little of the oxygen at this level is dissociated and thus little ozone is formed because almost all the high-energy UV has been filtered from sunlight before it descends to this altitude. For this reason, the ozone layer does not extend much below the stratosphere. Indeed, the ozone present in the lower stratosphere is largely formed at higher altitudes and over equatorial regions and is transported there. In contrast, at the top of the stratosphere, the UV-C intensity is greater, but the air is thin and therefore relatively little ozone is produced since the oxygen atoms collide and react with each other rather than with the small number of intact 02 molecules. Consequently, the production of ozone reaches a maximum where the product of UV-C intensity and 02 concentration is maximum. The maximum density of ozone occurs lower—at about 25 km over tropical areas, 21 km over mid-latitudes, and 18 km over subarctic regions—since much of it transported downward after its production. Collectively, most of the ozone is located in the region between 15 and 35 km, i.e., the lower and middle stratosphere, known informally as the ozone layer (see Figure 1-la).

A third molecule, which we will designate as M, such as N2 or HzO or even another 02 molecule, is required to carry away the heat energy generated in the collision between atomic oxygen and 02 that produces ozone. Thus the reaction above is written more realistically as

The release of heat by this reaction results in the temperature of the stratosphere as a whole being higher than the air that lies below or above it, as indicated in Figure 1-lb .

Notice from Figure 1-lb that within the stratosphere, the air at a given altitude is cooler than that which lies above it. The general name for this phenomenon is a temperature inversion. Because cool air is denser than hot air (ideal gas law), it does not rise spontaneously, due to the force of gravity; consequently, vertical mixing of air in the stratosphere is a very slow process compared to mixing in the troposphere. The air in this region therefore is stratified—hence the name stratosphere.

In contrast to the stratosphere, there is extensive vertical mixing of air within the troposphere. The Sun heats the ground, and hence the air in contact with it, much more than it does the air a few kilometers higher. It is for this reason that the air temperature falls with increasing altitude in the troposphere; the rate of decline of temperature with height is called the lapse rate. The less dense, hotter air rises from the surface and gives rise to extensive vertical exchange of air within the troposphere.


Given that tne total concentration of molecules in air decreases with increasing altitude, would you expect the relative concentration of ozone, on the ppb scale, to peak at a higher or a lower altitude or the same altitude compared to the peak for the absolute concentration of the gas?

Destruction of Stratospheric Ozone

The results for Problem 1-2 show that photons of light in the visible range and even in portions of the infrared range of sunlight possess sufficient energy to split an oxygen atom from a molecule of 03, However, such photons are not efficiently absorbed by ozone molecules; consequently, their dissociation by such light is not important, except in the lower stratosphere where little UV penetrates. As we have seen previously, ozone does efficiently absorb UV light with wavelengths shorter than 320 nm, and the excited state thereby produced does undergo a dissociation reaction. Thus absorption of a UV-C or UV-B photon by an ozone molecule in the stratosphere results in the decomposition of that molecule. This photochemical reaction accounts for much of the ozone destruction in the middle and upper stratosphere:

03 + UV photon (A < 320 nm)-> 02* + O*

The oxygen atoms produced in the reaction of ozone with UV light have an electron configuration that differs from the configuration that has the lowest energy, and they therefore exist in an electronically excited state; the oxygen molecules also are produced in an excited state.


By reference to the information in Problem 1-2, calculate the longest wavelength of light that decomposes ozone to O* and 02*, given the following thermochemical data:

02-» 02* AH0 = 95 kj moF1

[Hints: Express the overall reaction 0/O3 decomposition as a sum of simpler reaction5 for which AH° values are available, and combine their AH° values according to Hess' law, which states that AH° for an overall reaction is the sum of the AH° values for the simpler reactions that are added together.]

Most oxygen atoms produced in the stratosphere by photochemical decomposition of ozone or of 02 subsequently react with intact 02 molecules to re-form ozone. However, some of the oxygen atoms react instead with intact ozone molecules and in the process destroy them, since they are converted to 02:

In effect, the unbonded oxygen atom extracts one oxygen atom from the ozone molecule. This reaction is inherently inefficient since, although it is an exothermic reaction, its activation energy is 17 kj moP1, a sizable one for uv-c o, o atmospheric reactions to overcome. Consequen- + O

tly, few collisions between 03 and O occur with sufficient energy to result in reaction.

To summarize the processes, ozone in the stratosphere is constantly being formed, decomposed, and re-formed during daylight hours by a series of reactions that proceed simultaneously, though at very different rates depending upon altitude. Ozone is produced in the stratosphere because there is adequate UV-C from sunlight to dissociate some 02 molecules and thereby produce oxygen atoms, most of which collide with other 02 molecules and form ozone. The ozone gas filters UV-B and UV-C from sunlight but is destroyed temporarily by this process or by reaction with oxygen atoms. The average lifetime of an ozone molecule at an altitude of 30 kilometers is about half an hour; in the lower stratosphere, it lasts for months.

Ozone is not formed below the stratosphere due to a lack of the UV-C required to produce the O atoms necessary to'form 03, because this type of sunlight has been absorbed by 02 and 03 in the stratosphere. Above the stratosphere, oxygen atoms predominate and usually collide with other O atoms to eventually re-form 02 molecules.

The ozone production and destruction processes discussed above constitute the so-called Chapman mechanism (or cycle), shown in Figure 1-9. Recall that the series of simple reaction steps that document how an overall chemical process, such as ozone production and destruction, occurs at the molecular level is called a reaction mechanism.

Even in the ozone-layer portion of the stratosphere, 03 is not the gas of greatest abundance or even the dominant oxygen-containing species; its relative concentration never exceeds 10 ppm. Thus the term ozone layer is something of a misnomer. Nevertheless, this tiny concentration of ozone is sufficient to filter all the remaining UV-C and much of the UV-B from sunlight before it reaches the lower atmosphere. Perhaps the alternative name ozone screen is more appropriate than ozone layer.

As in the case of stratospheric ozone, it is not uncommon to find that the concentration of a substance, natural or synthetic, in some compartment of the environment or in an organism does not change much with time. This does not necessarily mean that there are no inputs or outputs of the substance. More often, the concentration does not vary much with time because the input rate and the rate at which the substance decays or is eliminated from some compartment in the environment have become equal; we say that the substance has achieved a steady state. Equilibrium is a special case of the steady state; it arises when the decay process is the exact opposite of the input. Box 1-1 explores the mathematical implications of the steady state in common situations involving reactive substances.


FIGURE 1-9 The

Chapman mechani sm.

The Steady-State Analysis of Atmospheric Reactions

BOX 1.1

The Steady-State Approximation

If we know the nature of the creation and destruction reaction steps for a reactive substance, we can often algebraically derive a useful equation for its steady-state concentration.

As a simple example, consider the formation and destruction of oxygen atoms above the stratosphere. As mentioned previously, the atoms are formed by the photochemical dissociation of molecules of diatomic oxygen:

The atoms re-form diatomic oxygen when two of them collide simultaneously with a third molecule, M, which can carry away most of the energy released by the newly formed 02 molecule:

Recall from introductory chemistry that the rates of the individual steps in reaction mechanisms can be calculated from the concentrations of the reactants and from the rate constant, k, for the step. Thus the rate of reaction (i) equals IcJOJ. The rate constant kt here incorporates the intensity of the light impinging upon the molecular oxygen. Since two O atoms are formed for each 02 molecule that dissociates, rate of formation of O atoms = 2 kJ02]

The rate of destruction of oxygen atoms by reaction (ii) is rate of destruction of O atoms = 2 kli[0]2[lS/I]

where we square the oxygen atom concentration because two of them are involved as reactants in the step.

The net rate of change of O atom concentration with time equals the rate of its formation minus the rate of its destruction:

rate of change of [O] = 2 kt[02] - 2 fcit[Of[M]

When atomic oxygen is at a steady state, this net rate must be zero, and thus the right-hand side of the equation above must also be zero. As a consequence, it follows that yo]2[M] = k,[o2]

By rearrangement of this equation, we obtain a relationship between the steady-state concentrations of O and of 02:

Continue reading here: [0ss2[02ss VCiiM

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  • Romola
    Which gas is responsible for absorption of UV in stratosphere in the range of 220320 nm?
    3 years ago
  • jyrki
    Why does ozone dissociate with radiation of a longer wavelength than oxygen?
    3 years ago
  • christina
    Why denser ozone layer undergoing rarification at some regions?
    4 years ago