4 -- CATALYTIC LOSS AND LIFE CYCLE OF A POLLUTANT

4.1 Chapman Chemistry

As alluded to earlier, Chapman chemistry does not give the full story behind the atmospheric ozone budget. Simple Chapman chemistry results in ozone amounts much greater than what are actually observed in the stratosphere. Figure 5.11 displays the total column ozone estimated using Chapman chemistry for March 21 conditions (solid line) with the actual satellite instrument measurements (dotted line). The measurements were made by the Total Ozone Mapping Spectrometer (TOMS) instrument. Details on the TOMS instrument are found in Chapter 7.

Chapman estimates are much too high in the tropics, and too low in the polar regions. The TOMS-derived global average for column ozone is about 300 Dobson Units (DU). The Chapman production and loss cycle produces a global average of about 790 DU. As noted already, there are two other mechanisms at work in addition to the Chapman cycle that reduce the equilibrium amount of ozone. First, reactions of OX with other trace gases alter the ozone budget. Second, the Brewer-Dobson Circulation transports ozone from its tropical source region into the extratropics. We will examine the effects of these other trace gases, such as hydrogen, nitrogen, chlorine, and bromine, on ozone in this section. The Brewer-Dobson Circulation is discussed in Chapter 6.

4.2 Catalytic Loss (Gas Phase)

Understanding catalytic loss processes is essential to understanding the observed distribution of stratospheric ozone. A catalyst is a substance, usually present in small amounts, that causes chemical reactions without itself being consumed by those reactions. Several catalytic cycles are important in ozone chemistry.

4.2.1 Hydrogen catalytic loss -- We first examine the impact of reactive hydrogen species (known as HOX, which consists of the hydroxyl radical, OH, and the hydroperoxy radical, HO2) on ozone chemistry. Hydrogen is transported into the stratosphere in the form of methane (CH4) and water vapor (H2O) molecules. Both of these molecules have large sources in the troposphere.

Methane, a long-lived tracer, is efficiently carried from its surface source region into the stratosphere by Brewer-Dobson tropical lifting motion. Methane is able to pass through the tropical tropopause into the lower stratosphere. Depletion of methane instead occurs in the upper stratosphere, where there is enough of the excited free oxygen atoms to trigger the oxidation reactions that converts methane into other species, including, water vapor (see Section 4.2.2).

Water vapor can reach very high amounts in the troposphere, as being caught outside on a rainy day shows, but is not efficiently transported into the stratosphere. Most of the water is "frozen out" as air moves through the very cold tropopause region in the tropical upper troposphere via the Brewer-Dobson Circulation. As moist air is lofted upward through the tropical tropopause, water vapor freezes into ice crystals, which do not readily enter the stratosphere. Water is carried into the stratosphere at only a few parts per million. Hence, the lower stratosphere tends to be very dry. Water vapor concentrations, however, increase in the middle to upper stratosphere and into the mesosphere due to the oxidation reactions of methane (see Section 4.2.2).

A typical methane distribution is shown in Figure 5.12 with data from the Halogen Occulation Experiment (HALOE) instrument aboard the Upper Atmospheric Research Satellite (UARS) taken March 5 to April 10, 1993. The figure shows that methane amounts are largest in the troposphere and decrease steadily through the stratosphere. The upward bowing in the tropics of methane mixing ratio contours is due to the lifting of the Brewer-Dobson Circulation. The source of methane in Earth's atmosphere can be traced to its release at the surface through a variety of sources: wood combustion, coal mining, oil and gas drilling and refining, landfills, wetland rice cultivation, crop residue burning, industrial activity, and digestive action by grazing animals (i.e., cow flatulence).

Figure 5.13 shows a typical water vapor distribution, again using data from HALOE taken during March and early April 1993. We see the same tropical lifting action where the water vapor mixing ratio contours are bowed upwards, but the amounts quickly fall off in the lower stratosphere. Unlike methane, though, water vapor actually begins to increase above the lower stratosphere, with elevated values in the mesosphere. The explanation for the difference in behavior of these two trace species is connected to both the chemistry and temperature structure of the atmosphere, which are presented below.

4.2.2 Methane photodissociation reactions -- Recall that the tropopause is the very cold boundary between the troposphere and the stratosphere. We've already seen that as moist air is lofted upward through the tropopause in the tropics, water vapor is frozen out. Air entering the stratosphere, consequently, is rather dry. On the other hand, methane remains unaffected by the cold temperatures as it passes through this boundary. Only when methane reaches the upper stratosphere is it depleted via oxidation reactions with OH (a very important player in HOX chemistry). These reactions lead to the production of water vapor molecules. Indeed, each methane molecule eventually is converted to (almost exactly) two molecules of water vapor in the middle to upper stratosphere via the following two oxidation reactions.

In the first reaction, methane is converted into water vapor by a reaction with the hydroxyl radical OH.

The second reaction involves a series of steps that begins with methane reacting with something called the singlet D oxygen atom, denoted O(1D). This is a free oxygen atom that is in a particular sort of excited state. The reaction is 

The result is a hydroxyl radical and a leftover methyl radical (CH3), which quickly reacts via

Eventually, reactions of HCHO (formaldehyde) with the hydroxyl radical result in the production of another water vapor molecule in the region between 35 and 45-km

Above about 65 km, photodissociation of methane becomes the dominate mechanism for methane loss (Le Texier, et al., 1988).

So the effect of the freeze drying of air at the cold tropical tropopause is to leave water vapor relatively scarce in the lower stratosphere, while oxidation of methane increases the relative amount of water vapor in the upper stratosphere. Methane, on the other hand, is relatively plentiful in the lower stratosphere (near its tropospheric source region), while it becomes relatively depleted in the upper stratosphere due to the above oxidation reactions. These processes help explain the appearance of water vapor and methane profiles through the stratosphere.

Figure 5.14 shows the sum of methane and water vapor (2*CH4+H2O). This figure demonstrates that despite the difference in the appearance of the vertical profiles of methane and water vapor, the two species are inversely related such that the total amount of hydrogen remains relatively constant throughout the stratosphere.

4.2.3 HOX catalytic cycles -- The importance of these two species, methane and water vapor, in ozone chemistry is that they transport and release hydrogen into the stratosphere. The activated hydrogen that is released can then participate in the destruction of odd oxygen, i.e., ozone, through a variety of catalytic cycles. These reactive hydrogen (HOX) cycles are summarized by Figure 5.15.

The open circles show the predominant species in which hydrogen exists in the stratosphere. We have not included H2 (molecular hydrogen), H2O (water vapor), and CH4 (methane) since they are not involved in the fast stratospheric hydrogen photochemistry balance. The arrows with superimposed boxes are reaction pathways. For example, OH (left circle) reacts with O3 to form HO2. The reaction is written

On the figure, we see this represented by the line with the superimposed blue (O, O3) box. The O2 (molecular oxygen) product is not represented, because it is not a hydrogen species. All of the reactions which lead to ozone creation are colored in blue, while ozone photolysis is colored in magenta.

Each water vapor molecule can be transformed into two molecules of HOX (reactive hydrogen) through reaction with O atoms via a reaction of water vapor with the singlet D oxygen atom.

Recall that HOX = OH + HO2. In this case, the reactive hydrogen exists in the form of two liberated OH (hydroxyl radical) molecules which become the catalyst in a pair of reactions with odd oxygen (OX) that result in a net loss of OX, by which we mean a net loss of both ozone molecules and free oxygen atoms.

This catalytic cycle involving HOx will only be disrupted if HOX is lost through another mechanism. Several reactions can remove HOx from this cycle

The hydrogen species on the righthand side of these reactions (i.e., H2O, HNO3, and HNO4) are known as reservoir species, which are chemical compounds that store (like a reservoir) a particularly species in a nonreactive form. These species act as stores of hydrogen, locking up or sequestering HOx and preventing its participation in the catalytic cycle outlined above. H2O, HNO3 and HNO4 react very slowly with odd oxygen. When HOx is tied up in one of the reservoir species, therefore, it is not important in the loss of odd oxygen.

Since reservoir species are relatively inert (compared to reactive species like OH and HO2), they tend to have very long lifetimes, and the less reactive they are, the longer are their lifetimes. Eventually, reservoir species are transported into the troposphere, removing HOx from the stratosphere altogether. H2O can be converted back into reactive hydrogen via the reaction with O(¹D) outlined above. HNO3 (nitric acid), on the other hand, under normal conditions in the stratosphere, releases its HOx mainly through photolysis. The photolysis of the nitric acid reservoir species is discussed in the section on NOx catalytic cycles.

In the upper stratosphere, where O atoms are relatively plentiful, the HOx catalytic cycle outlined above (and highlighted in Figure 5.15a) is fairly effective. In the lower stratosphere, however, a different catalytic cycle involving HOx is responsible for Ox loss.

Again, notice that HOx (OH + HO2) is neither produced nor destroyed by this cycle, but merely acts as a catalyst for converting two molecules of ozone into three molecules of O2. The importance of this reaction is that it does not require free O atoms for the reaction

Two other reactions are important for OX loss in the upper stratosphere. The first one involves a free H (hydrogen) atom as an intermediate compound. See Figure 5.15c for the pathway diagram. Two O atoms are converted into a single diatomic oxygen molecule.

The second reaction involves HOx and the loss of two odd oxygens (an ozone molecule and an oxygen atom). See Figure 5.15d for the pathway diagram. The two odd oxygens are converted into two diatomic oxygen molecules

Finally, yet another pair of catalytic cycles is important in the lower stratosphere. These cycles involve interaction with the chlorine or bromine cycles (addressed below). These cycles are the most complicated of the catalytic cycles we've looked at so far, but yield the same result: catalytic destruction of odd oxygen. In these reactions, Z can be either chlorine (Cl) or bromine (Br):

The two cycles are summarized in the pathway diagram in Figure 5.15e.

4.2.4 NOX catalytic cycles -- The loss of odd oxygen can also occur through catalytic cycles involving nitrogen species in the form of reactive nitrogen. Reactive nitrogen, denoted NOx, includes NO (nitric oxide) and NO2 (nitrogen dioxide). Like reactive hydrogen, reactive nitrogen species have their origins in the troposphere. Approximately 90% of stratospheric NOx comes from tropospheric N2O (nitrous oxide, also known as laughing gas). Like H2O and CH4, N2O is transported upward into the stratosphere mainly through the tropical tropopause by the lifting of the Brewer-Dobson Circulation. Figure 5.16 shows a typical distribution of N2O in the stratosphere as seen by the Cryogenic Limb Array Etalon Spectrometer (CLAES) instrument aboard the Upper Atmospheric Research Satellite (UARS) in January 1993. This corresponds to wintertime in the northern hemisphere. Recall that the Brewer-Dobson Circulation exists in the winter hemisphere between equator and pole. Notice in Figure 5.16 that the highest values of N2O are observed near its source region in the lower, tropical troposphere.

Reactive nitrogen species are formed from N2O via the reaction of nitrogen dioxide with the singlet D oxygen atom to form two molecules of nitric acid.

This reaction transfers nitrogen from the inert species, N2O, into the reactive species, NO.

Another way that N2O is lost is via photolysis. An energetic UV photon is able to dissociate N2O into molecular nitrogen, N2, and the singlet D oxygen atom.

Since N2 is a very long-lived, nonreactive species, it does not contribute to photochemical processes, even though is the most abundant gas in the atmosphere.

Sources of N2O at the ground include oceans, forest soils, combustion, biomass burning, and fertilizers. The amount of nitrogen generated annually by these processes are estimated to be in the 4.4 to 10.5 teragrams (Tg). (One Tg=10¹² or one trillion grams.) Some of the chemical processes of these nitrogen oxides are illustrated in Figure 5.17, which shows the NOX catalytic cycles.

The circles show many of the nitrogen containing species in the stratosphere. We have not included N2, since it is not involved in stratospheric nitrogen photochemistry, nor have we included N2O, since it is the prime source of NO. The reaction pathways are shown as dark arrows with superimposed boxes. For example, BrONO2 (bromine nitrate, shown in top circle) can be photolyzed by energetic photons to form Br (bromine) and NO3 (nitrogen trioxide). The reaction is given by

Br is not represented, because it is not a nitrogen species. All of the reactions which lead to ozone loss (i.e., OX loss) are shown by the dark arrows with the superimposed blue boxes. Photolysis reactions are shown by the dark arrows with superimposed magenta boxes inside of which is written h. We can represent the energy of a photon by either hc/ or h, since the relationship between the speed of light (c), wavelength (), and frequency () is given by

The reactive forms of nitrogen drive their own catalytic cycle, analogous to the HOX cycles. The reactions are represented in Figure 5.18 by the red and yellow highlights. In step A (lower yellow line), NO reacts with O3 (indicated by the red O3 in the blue box on the lower yellow line) to form NO2 and O2. The NO2 then reacts with O to reform NO (upper yellow line). In the process, both an ozone and oxygen molecule are destroyed (i.e., two OX molecules), while the NO is reformed

The catalytic cycle results in the loss of two odd oxygen without loss of the catalytic NOX species.

4.2.5 Temperature dependence of NOX catalytic reactions -- Reactions between chemical species are temperature dependent. This temperature dependence for NOx species is particularly strong. Figure 5.19, left side, shows the reaction rate of NO+O3 --> NO2+O2 (expressed in units of cm³/sec/molecule) as a function of temperature over a range of temperatures characteristic of the stratosphere. This type of reaction is referred to as a bimolecular reaction, as it involves two molecules. Note that the NO+O3 rate increases very steeply as the temperature increases. This steep increase is associated with the faster speeds and energies that molecules acquire at higher temperatures. From the statistical behavior of large collections of molecules, we know the average energy of a molecule is related to its temperature. The warmer the temperature (i.e., the faster the molecules are moving), the greater the probability that two molecules will meet and react. Hence, the reaction rate increases with temperature. These reactions rates are also dependent on other factors, such as the energy associated with the molecular bonds.

In contrast to these bimolecular (two molecule) reactions, the termolecular (three molecule) reaction rate tends to decrease with increasing temperature. The right panel of Figure 5.19 shows the reaction rate of O+O2+M-->O3+M (expressed in units of cm³/sec/molecule) as a function of temperature over a range of temperatures characteristic of a 2 mb pressure (about 40 km in altitude). In this reaction, the rate decreases as temperature increases. The reason for this decrease is that the probability of three molecules colliding simultaneously to react drops sharply as the molecules are moving about faster (i.e., as the temperature rises).

The temperature dependence of these NOx reaction rates can be used to show why ozone and temperatures are anticorrelated in the upper stratosphere. That is, we can show why ozone concentrations fall as temperature rises and vice versa. Recall that the change in concentration of a molecule is due to both photolysis and chemical reactions. The following three chemical equations allow us to solve for ozone concentration as a function of temperature and NOx. 

(1) Change in OX due to NOx catalytic loss and photochemical production

Our first equation arises from the fact that the change of Ox (i.e., ozone and free oxygen) is due to a balance between NOx catalytic loss and photochemical production. Mathematically, we may write this as

The first term on the left is ozone loss by the NO + O3 reaction, the second term is O loss by the NO2 + O reaction, the 1st term on the right is the photolysis of NO2 by solar radiation which produces an O atom and an NO molecule, while the last term is the photolysis of O2 which produces two oxygen atoms. (We've neglected other processes here for illustration purposes.)

(2) Swapping back and forth of NOX between NO and NO2

The second equation arises from the fact that NOX is constantly being swapped back and forth from NO to NO2, which we can mathematically express as

This shows us mathematically that the production of NO2 is equal to the loss of NO2. The term on the left hand side of the equation is NO2 production by the NO+O3 reaction. The first term on the right hand side is the loss of NO2 by the NO2+O reaction, and the second term is the loss of NO2 by photolysis.

(3) Swapping back and forth of Ox between O and O3

The third equation arises because odd oxygen, Ox, like reactive nitrogen, NOx, is constantly being swapped back and forth from free oxygen, O, to ozone, O3. We can mathematically express this as

The first term is the ozone production by the O+O2+M termolecular reaction, while the term on the right hand side is the photolysis of ozone.

We can combine the three equations above to solve for the ozone concentration as a function of temperature and NOx. This is a relatively complicated expression. However, Figure 5.20 shows the graphical results of such a calculation by making a few assumptions. We first assume a NOx concentration of 2.0 x109 molecules/cm³. We then perform the calculation at the Equator on March 21, when the Sun is directly overhead. We then assume a pressure level of 2 mb (about 40 km altitude). We then are able to derive an ozone-versus-temperature plot where ozone is a function of temperature.

We see in Figure 5.20 that ozone amounts decrease with increasing temperature. The decrease is a result of the temperature dependence of the NO+O3 reaction which is strongly temperature dependent (as shown in the left panel of Figure 5.19). As temperatures increase, this reaction speeds up, destroying ozone faster, and reducing the ambient amount of ozone. This reaction rate temperature dependence explains the observed ozone-temperature relationship in the upper stratosphere, and illustrates the impact of temperature dependent reaction rates on the stratosphere.

4.2.6 Interference cycles -- As with the hydrogen atoms cycling back and forth between OH and HO2, nitrogen atoms can take several pathways as they cycle between NO and NO2. Those pathways in which an odd oxygen molecule is not lost are referred to as interference cycles. Figure 5.21 illustrates an interference cycle. The reactions involved are

In this cycle, NOx does not act as a catalyst to destroy ozone, even though it reacts with ozone in the second step. The first reaction (photolysis of NO2) creates an O atom (i.e. production of odd oxygen), while the second step destroys an O3 molecule (i.e., loss of odd oxygen). The third reaction recreates the ozone molecule, leading to no net change. This reaction chain is relatively effective in interfering with the normal nitrogen catalytic loss.

There are several termolecular reactions which can transfer nitrogen from reactive forms (NO and NO2) into less reactive forms

The species on the righthand side of the equations are known as reservoir species and are relatively nonreactive with odd oxygen species. Each of these nitrogen reservoirs can release NOx through photolysis. Their photolysis rates are quite different, however, and characterize the lifetime of the various reservoir species. In general, HNO3 (nitric acid) has the longest lifetime while N2O5 (dinitrogen pentoxide) has the shortest.

As was the case with the lifetimes of O2 and O3, the lifetime of the reservoir species is controlled by the photolysis rate. Since these rates depend upon the intensity of the incoming solar radiation, they will vary with time of day, latitude, altitude, and season.

Let's look at the photolysis rates under the following conditions: a spring day in Washington, DC, near noon at an altitude of 30 km. Under such conditions, the maximum photolysis rate of N2O5 is around 9x10^-5 sec^-1 (corresponding to an e-folding time of 3 hours), of ClONO2 (chlorine nitrate) around 1x10^-4 sec^-1 (corresponding to an e-folding time of 2.6 hours), and of HNO3 around 1.4x10^-5 sec^-1 (corresponding to an e-folding time of 20 hours). The difference in these reaction rates suggest that N2O5 and ClONO2 should vary noticeably throughout the course of the day, with substantially less present in the late afternoon than in the early morning. The slower photolysis of HNO3 molecules suggest that they are more effective reservoirs of NOx, requiring more hours than available daylight to be transformed from reservoir species back into active species. Most of these molecules will survive the course of a day. Variation in HNO3 will be observed only over much longer timescales. At 20 km, HNO3 is an even more effective reservoir, with a photolysis rate of 10^-6 (corresponding to an e-folding time of 10 days!). Once freed by photolysis, NOx can return to participate in the catalytic cycles outlined above. The effect of locking up nitrogen into reservoir species like nitric acid is that the nitrogen can't destroy ozone! 

4.2.7 Chlorine sources -- As with the nitrogen and hydrogen species, the sources of chlorine are also in the troposphere. Chlorine atoms are bound up in the various manmade chlorofluorocarbon (CFC) and hydrochlorofluorocarbon (HCFC) molecules, including F-11, f12, f113, f114, CCl4, CH3CCl3, CH3Cl, HCFC-22. These molecules can be transported across the tropopause into the stratosphere. CFCs were developed in the 1920s as a safe, nontoxic refrigerant alternative to ammonia (see Chapters 1 and 11). That is, both CFCs and ammonia make good coolants. However, when CFCs leak, there are no adverse health consequences, while when ammonia leaks, undesirable health consequences, even death, can occur.

In addition to their characteristic as a good refrigerant, CFCs are cheap to manufacture, nonflammable, and insoluble (meaning that when released in the atmosphere, they would not be captured in rain drops and deposited on Earth's surface). They gained enormous usage worldwide from the 1940s to the 1980s.

Because CFCs have very long lifetimes and are not water soluble, they can be transported upward into the stratosphere. Indeed, the current stratospheric chlorine content arises mostly from CFCs. This represents a dramatic change to the atmosphere by human activity. While natural sources of chlorine exist, it is the release of chlorine atoms from the photolysis of CFC molecules that provide most of the observed chlorine in the stratosphere at the present time. In the upper stratosphere, the high energy UV radiation above the ozone layer can break the CFC bonds, and releasing chlorine to participate in its own catalytic cycle destroying odd oxygen. We will explore these catalytic cycles in Section 4.2.8. 

4.2.8 ClX catalytic reactions -- The chemical reactions of different chlorine-containing molecules are illustrated in Figure 5.22. The circles show the predominant short-lived species that form the chlorine containing family in the stratosphere. We have not included CFCs on the graph because of there relatively long lifetimes. The reaction pathways are drawn as black arrows with superimposed boxes. For example, Cl (left, middle circle) can react with O3 to form ClO (chlorine monoxide) and O2. The reaction is written 

On the figure, we see this represented by the line with the O3 blue box superimposed on the line. The O2 is not represented, because it is not a chlorine species. All of the reactions which lead to ozone loss (i.e., OX loss) are colored in blue, while photolysis reactions are colored in magenta.

A principal loss of ozone in the upper stratosphere is the Cl/ClO reaction, represented in Figure 5.23, and written as

Here the reactive chlorine compounds are Cl and ClO, denoted ClX. The net reaction is to convert two odd oxygens, an ozone molecule and an oxygen atom, into two molecules of diatomic oxygen. Since there is little free oxygen in the lower stratosphere, this reaction is not the principal loss mechanism for polar lower stratospheric ozone.

As in NOX case, several reactions exist that transform reactive chlorine into reservoir species.

The chlorine reservoir species HCl (hydrochloric acid), HOCl (hypochlorous acid), and ClONO2 (chlorine nitrate) are characterized by a variety of lifetimes, determined by their photolysis rates. HCl is the longest lived, with a lifetime on the order of weeks. HOCl is the shortest, with a lifetime on the order of hours, as HOCl quickly photolyzes in sunlight.

In the lower stratosphere, several other catalytic cycles involving chlorine have important effects on the ozone balance. Yet another cycle involves the photolysis of NO3 and ClONO2. This cycle is highlighted in Figure 5.24. The reactions are

The three-body reaction of ClO with NO2 and some other molecule M deactivates the reactive chlorine species ClO into nonreactive chlorine species ClONO2. The net effect of this set of reactions involving both reactive and nonreactive forms of chlorine and nitrogen is to convert two molecules of ozone into three molecules of diatomic oxygen.

4.2.9 ClX catalytic reactions and the Antarctic ozone hole -- Chapter 11 details the chemical reactions responsible for destroying stratospheric ozone in the Antarctic each spring, leading to the formation of the "ozone hole." Here we will consider two critical Clx catalytic cycles in the ozone hole phenomenon. In both cases, two ozone molecules are converted into three diatomic oxygen molecules.

(1) First catalytic cycle

(2) Second catalytic cycle (shown in Figure 5.25):

Such a sequence of reactions effectively transforms one form of Ox to another with no net loss, while preventing NOx and Clx molecules from participating in normal catalytic cycles. Such interference cycles effectively slow destruction of Ox.

4.2.10 Bromine sources -- Bromine is another molecule effective for ozone destruction. Bromine accounts for significant ozone loss (about 20-40%) inside the Antarctic ozone hole. The sources of bromine are both anthropogenic and natural. Methyl bromide (CH3Br) is produced in the troposphere, but it is also the predominant source of bromine in the stratosphere. Methyl bromide is produced by biological processes on both land and in the ocean. Methyl bromide is also used as a fumigant for agricultural purposes, and is released via biomass burning and from cars using leaded fuel. Two other major sources of bromine are Halon-1211 (CBrClF2) and Halon-1301 (CBrF3), used as fire suppressants.

The losses of methyl bromide occur through reactions with water, the hydroxyl radical, chlorine ions, and photolysis by ultraviolet radiation. There are probably also losses via biological processes, but these are uncertain. Methyl bromide has a long lifetime, which allows some of it to be lifted out of the troposphere and into the stratosphere.

Halons 1211 and 1301 are only destroyed by UV photolysis at wavelengths shorter than 280 nm. Hence, the halons can only be photolyzed in both the upper and lower stratosphere. They have very long lifetimes, since it takes quite a while for a molecule to reach these altitudes. The lifting action is again provided by the Brewer-Dobson Circulation.

4.2.11 BrX catalytic reactions -- Reactive bromine exists in the form of bromine (Br) and bromine monoxide (BrO). Nonreactive bromine species include hypobromous acid (HOBr) and bromine nitrate (BrONO2). These are typically not referred to as "reservoir species" because they are very easily photolyzed, even by visible light, and hence have very short lifetimes. This means that they do not lock up reactive bromine in the same way that ClONO2 locks up reactive chlorine, and so bromine species in the stratosphere tend to exist in reactive forms.

The chemical processes for bromine are illustrated in Figure 5.26. The open circles show these main bromine species in the stratosphere, all of which are relatively short-lived. We have not included methyl bromide or any halons, since their lifetimes are long. The reaction pathways are shown as black arrows with superimposed boxes. One such pathway involves the reaction of Br (left, bottom circle) with ozone (O3) to form BrO and O2. The reaction is written

On Figure 5.26, we see this represented by the black arrow with the superimposed O3 blue box. O2 is not represented, because it is not a bromine species. As in the previous catalytic cycles for reactive hydrogen, nitrogen, and chlorine species, the bromine catalytic reactions which lead to ozone (i.e., OX loss) are colored in blue, while photolysis reactions of bromine species are colored in magenta.

There are four distinct catalytic cycles for ozone loss. These reaction cycles are illustrated in Figure 5.27 (A, B, C, and D). The principle reactants are highlighted in red. The reaction pathways are shown by the yellow arrows. The yellow shaded circles show the principle bromine reactants.

(A) Brx - Ox Reaction Cycle -- In this cycle, a two-step reaction occurs between reactive bromine and odd oxygen. First, a BrO molecule reacts with a free O atom to form a free Br atom and a diatomic oxygen molecule. Next, the Br atom reacts with an O3 molecule to reform the BrO molecule and a molecule of diatomic oxygen. The net reaction is to convert two odd oxygen species, an O atom and an O3 molecule, into two diatomic oxygen molecules. The reactions are

(B) Brx - Clx - Ox Reaction Cycle -- In this second cycle, four reactions occur involving reactive Brx and Clx species and odd oxygen, in the form of ozone. The net result is to convert two O3 molecules into three O2 molecules. That is, two odd oxygen species are lost. The reactions are

(D) Brx - HOx - Ox Reaction Cycle -- In this fourth cycle, there are four reactions among reactive BrX and HOx species and Ox, again in the form of O3. Also involved is another nonreactive form of bromine (HOBr), which is also quickly photolyzed by near-UV and visible light. The net change is as in (B) and (C) above: the conversion of two O3 molecules into three O2 molecules, representing the loss of Ox. These reactions are

These chains of reactions make bromine one of the most efficient destroyers of ozone for two reasons. First, catalytic cycles (B), (C), and (D) do not require free oxygen atoms to destroy ozone, meaning that these reactions can occur in the lower stratosphere where there are few O atoms available. Second, HOBr and BrONO2 are very easily photolyzed, so that bromine typically exists as reactive species (Br or BrO) rather than as nonreactive species (BrONO2 and HOBr).