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.
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Those pathways in which an odd oxygen molecule is not lost are referred to as interference cycles. The reactions involved are. In this cycle, NO x does not act as a catalyst to destroy ozone, even though it reacts with ozone in the second step. The first reaction photolysis of NO 2 creates an O atom i. 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 NO 2 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 NO x through photolysis. Their photolysis rates are quite different, however, and characterize the lifetime of the various reservoir species. In general, HNO 3 nitric acid has the longest lifetime while N 2 O 5 dinitrogen pentoxide has the shortest. As was the case with the lifetimes of O 2 and O 3 , 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 N 2 O 5 is around 9x10 - 5 sec - 1 corresponding to an e-folding time of 3 hours , of ClONO 2 chlorine nitrate around 1x10 - 4 sec - 1 corresponding to an e-folding time of 2.
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The difference in these reaction rates suggest that N 2 O 5 and ClONO 2 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 HNO 3 molecules suggest that they are more effective reservoirs of NO x , 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 HNO 3 will be observed only over much longer timescales.
At 20 km, HNO 3 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, NO x 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! These molecules can be transported across the tropopause into the stratosphere.
CFCs were developed in the s as a safe, nontoxic refrigerant alternative to ammonia see Chapters 1 and 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 s to the s. 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. 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. On the figure, we see this represented by the line with the O 3 blue box superimposed on the line. The O 2 is not represented, because it is not a chlorine species. 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 NO X case, several reactions exist that transform reactive chlorine into reservoir species. The chlorine reservoir species HCl hydrochloric acid , HOCl hypochlorous acid , and ClONO 2 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. In the lower stratosphere, several other catalytic cycles involving chlorine have important effects on the ozone balance. This cycle is highlighted in Figure 5. The reactions are. 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.
In both cases, two ozone molecules are converted into three diatomic oxygen molecules. There are no O atoms involved in either of these two catalytic cycles. Yet reactive chlorine is still created when the ClO dimer Cl 2 O 2 is photolyzed by UV light, liberating chlorine atoms, which then destroy ozone.
Hence, these reactions can destroy ozone in the lower stratosphere, where there are very few O atoms. However, it turns out that these reactions are relatively unimportant throughout most of the stratosphere most of the time. It is only in the presence of polar stratospheric clouds PSCs , which are described in Chapter 11, that chlorine reservoir species are liberated. In addition, it is only in the presence of PSCs that reactive nitrogen species, NO x , are locked up as nitric acid, which stops the reaction of NO 2 and the reactive chlorine species ClO.
This allows ClO levels to increase dramatically. As discussed in the case of nitrogen, interference cycles can reduce ozone loss rates. For example,. Such a sequence of reactions effectively transforms one form of O x to another with no net loss, while preventing NO x and Cl x molecules from participating in normal catalytic cycles. Such interference cycles effectively slow destruction of O x. The sources of bromine are both anthropogenic and natural. Methyl bromide CH 3 Br 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. 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 and are only destroyed by UV photolysis at wavelengths shorter than 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. 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 ClONO 2 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. These studies have shown that several reactions involving chlorine and bromine directly or indirectly destroy ozone in the stratosphere. Computer models have been used to examine the combined effect of the large group of known reactions that occur in the stratosphere.
These models simulate the stratosphere by including representative chemical abundances, winds, air temperatures, and the daily and seasonal changes in sunlight. These analyses show that under certain conditions chlorine and bromine react in catalytic cycles in which one chlorine or bromine atom destroys many thousands of ozone molecules. Models are also used to simulate ozone amounts observed in previous years as a strong test of our understanding of atmospheric processes and to evaluate the importance of new reactions found in laboratory studies.
The responses of ozone to possible future changes in the abundances of trace gases, temperatures, and other atmospheric parameters have been extensively explored with specialized computer models.wordpress-11600-25562-61096.cloudwaysapps.com/14703.php
Basic Ozone Layer Science | Ozone Layer Protection | US EPA
Atmospheric observations have shown what gases are present in different regions of the stratosphere and how their abundances vary. Gas and particle abundances have been monitored over time periods spanning a daily cycle to decades. Observations show that halogen source gases and reactive halogen gases are present in the stratosphere at the amounts required to cause observed ozone depletion.
Ozone and chlorine monoxide ClO , for example, have been observed extensively with a variety of instruments.
Ozone reaction mechanisms
ClO is a highly reactive gas that is involved in catalytic ozone destruction cycles throughout the stratosphere. Instruments on the ground and on satellites, balloons, and aircraft now routinely detect ozone and ClO remotely using optical and microwave signals. High-altitude aircraft and balloon instruments are also used to detect both gases locally in the stratosphere. The observations of ozone and reactive gases made in past decades are used extensively in comparisons with computer models in order to increase confidence in our understanding of stratospheric ozone depletion.
Chemicals that destroy ozone are formed by industrial and natural processes. With the exception of volcanic injection and aircraft exhaust, these chemicals are carried up into the stratosphere by strong upward-moving air currents in the tropics. Methane CH 4 , chlorofluorocarbons CFCs , nitrous oxide N 2 O and water are injected into the stratosphere through towering tropical cumulus clouds. These compounds are broken down by the ultraviolet radiation in the stratosphere. Aerosols and clouds can accelerate ozone loss through reactions on cloud surfaces.
Formation of ozone
Thus, volcanic clouds and polar stratospheric clouds can indirectly contribute to ozone loss. Stratospheric air temperatures in both polar regions reach minimum values in the lower stratosphere in the winter season. This occurs on average for 1 to 2 months over the Arctic and 5 to 6 months over Antarctica see heavy red and blue lines. Reactions on PSCs cause the highly reactive chlorine gas ClO to be formed, which increases the destruction of ozone.