The stratosphere is situated between the troposphere and the mesosphere, at an altitude of roughly 12-50km, although this is variable. It tends to start lower in the polar regions, and higher nearer the equator. It is a region in the atmosphere where the temperature rises with altitude. However, due to a lack of convection, substances diffuse slowly, so reactions here tend to be altitude dependent. Ozone destruction happens naturally in the atmosphere via the Chapman Mechanism, which is dramatically increased by chloride and bromide radicals8.
Thomas Midgley first discovered Chlorofluorocarbons (CFCs), dubbed Freon, in 1928. They are compounds which contain carbon, fluorine and chlorine atoms, and were used in refrigeration, aerosols, bedding, as pesticides and had many other uses. Hydrochlorofluorocarbons (HCFCs) are CFCs which also contain hydrogen atoms. They have similar properties, and therefore uses, so were used as temporary CFC replacements, along with Hydrofluorocarbons (HFCs) when CFC use was banned by the Montreal Protocol in 1987. HFCs are also similar, but do not contain any chlorine. Principle CFCs used include CCl3F and CCl2F2, and a principle HCFC is CHClF2. There are also many bromine containing compounds emitted into the atmosphere; CFCs containing bromine are known as halons, which are used as fire suppressants, examples of which are CF2ClBr and CF3Br. The most abundant bromine containing compound emitted is methyl bromide (CH3Br), which comes from both natural and manmade sources. CH3Br makes up 55-70% of bromine in the stratosphere contributed from CH3Br, CH2Br2 and the two halons.
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In the stratosphere, CFCs absorb UV radiation and photodissociate, producing chlorine radicals.
CFCl3 + hν à CFCl2 + Cl
F + CF3H à CF3 + HF (1)
FO2 + O3 à FO + 2O2 (slow reaction) (2)
FO + O3 à FO2 + O2 (3)
The maximum photolysis rate of CFCl3 is ~25km, CF2Cl2 ~ 32km and CClF2CF3 ~ 40km8. The more heavily chlorinated molecules photolyse at lower altitudes, and so have a greater atmospheric impact. HCFCs are reactive in the troposphere, unlike CFCs (which are inert here), and so can be destroyed by OH radicals. Those which make it to the stratosphere undergo photolysis like CFCs. HFCs can also be destroyed in the troposphere by the same mechanism, but do not contain chlorine, so cannot photolyse to produce chlorine radicals. As the C-Br bond is weaker than the C-Cl bond, bromine compounds photolyse lower in the stratosphere,. The absorption cross sections of bromine compounds also tend to be larger than those of corresponding chlorine compounds, especially at longer wavelengths, where there is a larger intensity of solar radiation available for photolysis8.
CBr3F + hν à Br + CBr2F
Cl + O3 à ClO + O2 (8)
ClO + O à Cl + O2 (9)
Net: O3 + Oà 2O2 (10)
CFCl2 + O2 + M à CFCl2O2 + M (4)
CFCl2O2 + NO à CFCl2O + NO2 (5)
CFCl2O + M à COFCl + Cl + M (6)
COFCl + hν à FCO + Cl (7)
The C-F has a higher bond dissociation energy than the C-Cl/Br bonds, so is harder to break, meaning photodissociation producing fluorine rarely happens. When produced, fluorine radicals can then react with HFCs (reaction 1)6. Fluorine radicals can also react with oxygen, then go on to destroy ozone (reactions 2-3)6. Hydrogen fluoride is unreactive in the stratosphere, so fluorine radicals can be removed by reacting with methane or water, which produces this. Ozone depletion by chlorine is more than 104 times more efficient than by fluorine6, and so the net effect of fluorine atom chemistry on ozone destruction is very small.
ClO + O à Cl + O2 (13)
ClO + HO2 à HOCl + O2 (14)
HOCl + hν à Cl + OH (15)
ClO + NO à Cl + NO2 (16)
NO2 + O à NO + O2 (17)
ClO + NO2 + M à ClONO2 + M (18)
ClONO2 + hν à Cl + NO3 (19)
NO3 + hν à NO + O2 (20)
Br + HCHO à HBr + CHO (11)
HBr + OH à Br + H2O (12)
Once photolysed, if there are chlorine atoms remaining on the CFC radical, it can react to produce further chlorine radicals (reactions 4-7). In these reactions M is a third body, needed to either give the molecule energy to put it into an excited state, or to remove excess energy from the reaction.
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Once produced, chlorine radicals can take place in a catalytic chain reaction with ozone, as can bromine (reactions 8-10) . This chemistry is primarily important in the middle and upper stratosphere. Unlike chlorine, atomic bromine doesn't react rapidly with organic molecules, but it does have a reasonable reaction rate with HCHO (reaction 11)6, which competes with the bromine version of reaction 8. HBr is converted back to atomic bromine via reaction 126.
In the lower stratosphere, the main destruction of ozone through halogens is through the HalOx cycles (where Hal is Cl or Br, reactions 13-20)8, the HalO being formed in reaction 8. This is due to the fact that they regenerate the atomic radicals as catalysts for ozone destruction.
BrO + NO à Br + NO2 (24)
BrO + HO2 à HOBr + O2 (25)
HOBr + O(3P) à OH + BrO (26)
ClO + BrO à Br + OClO (21)
àBr + ClOO (22)
àBrCl + O2 (23)
The ClOx and BrOx cycles are connected by reaction with each other (reactions 21-23)6. Although the reaction pathway producing BrCl is only a minor channel, BrCl absorbs in the UV/Vis region, so it ultimately regenerates atomic bromine and chlorine. Reactions 13-16 show a pathway which accounts for ~30% of ozone loss by halogens in the lower stratosphere for chlorine, and the bromine equivalent ~20-30%6.They are also significant in lower polar stratosphere. BrO also reacts to give atomic bromine with itself (generating O2) and NO (reaction 24)6, and with O(3P) in the upper stratosphere.
BrO is the major form of bromine in the lower stratosphere, despite the rapid sink processes (reaction 25), and the fact that the lifetime of BrO with respect to photolysis in the stratosphere is of the order of seconds.
Above 25km, the reaction of HOBr is important (reaction 26)6. HOBr also indirectly couples chlorine and bromine chemistry, as its photolysis increases the concentration of OH radicals, which causes faster recycling of HCl back to atomic chlorine. ClO reacts with itself in the same way that it does with bromine (reactions 21-23), although they occur at much slower rates (in the order of 10-15 as opposed to 10-12), and so are not as significant.
Cl atoms are also reformed through the thermal decomposition of ClOO to Cl + O2 (via infrared), and of OClO to give O + ClO. The second of these reactions doesn't lead to a net loss of ozone, as the oxygen atom can react with molecular oxygen to regenerate ozone6.
H2O + ClONO2 (particle) à HOCl + HNO3 (33)
HOCl + HCl (particle) à Cl2 + H2O (34)
Net: HCl(ads) + ClONO2 à Cl2 + HNO3 (35)
HCl(ads) + N2O5 à ClNO2 + HNO3 (36)
Br + HO2 à HBr + O2 (31)
BrONO2 + hν à BrO + NO2 (32a)
à Br + NO3 (32b)
Cl + CH4 à HCl + CH3 (27)
ClO + NO2 + M à ClONO2 (28)
HCl + OH à Cl + H2O (29)
ClONO2 + hν à Cl + NO3 (30)
Both chlorine and bromine radicals can be removed from the stratosphere into temporary reservoir compounds. Chlorine radicals can react with methane or nitrogen dioxide (reactions 27-28)8, forming the reservoir compounds HCl and ClONO2, and is regenerated through the reactions 29-30. Bromine radicals don't react fast with methane, instead forming HBr through reaction with HO2 (reaction 31)8. It also reacts is the same way as chlorine in reaction 28. However, these reservoirs of bromine are not very efficient, and photolyse much more rapidly than the chlorine equivalents (reactions 22a-b)8. This means that bromine spends a lot more time in its catalytically active form, and so is a lot more destructive to ozone.
ClO + ClO + M à (ClO)2 (37)
(ClO)2 + hν à Cl + ClOO (38)
ClOOà Cl + O2 (39)
In the poles, reactions can take place on the surface of polar stratospheric cloud particles (PSCs), which allow some reactions to take place, and others to occur many times faster than in gaseous phase (reactions 33-36)6. Reactions 37-396 take place in the ozone hole in the Antarctic, and this is the most important channel here. The ozone hole forms due to the build up of chlorine reservoirs during the winter, in which there is no sunlight, in the Antarctic vortex. In spring, with the return of sunlight, there are lot more reservoir compounds to photolyse, and so there is a sudden decrease in ozone concentration as the radicals produced destroy it (Figure 1). The difference between winter (August), and spring (October), is very obvious. The arctic vortex is a lot weaker, and so there isn't an ozone hole fixed there, it moves around.
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Figure 1 Partial pressures of ozone in nanobars versus altitude, as measured by balloon sondes from McMurdo, Antarctica.
In conclusion, bromine containing compounds in the stratosphere are the most destructive towards ozone, as it spends longer in a catalytically active form than chlorine, even though they both have very similar reaction pathways. CFCs are inert in the troposphere, like the compounds containing bromine, so are much more destructive than HCFCs, which are not, as the produce a lot more chlorine radicals in the stratosphere. HFCs are the least destructive towards stratospheric ozone, as they can be destroyed in the troposphere, and do not have much ozone destructive chemistry as they do not contain any chlorine.
The troposphere is the first layer of the atmosphere, extending up to 8km in the poles and 16km at the equator. It contains primarily nitrogen and oxygen, with only small trace gas concentrations. Most weather takes place here. There are many different sulphur containing compounds present in the troposphere, emitted from both natural and anthropological sources. The vast majority of sulphur emitted is in the form of SO2, with small amounts of H2SO4, sulphates and SO3. The other main inorganic sulphur compounds present are hydrogen sulphide (H2S), carbon disulfide (CS2) and carbonyl sulphide (COS), and the organics are dimethyl sulphide (CH3SCH3, DMS), dimethyl disulfide (CH3SSCH3, DMDS), and methyl mercaptan (CH3SH).
The three inorganic compounds all react with the hydroxyl radical to produce the SH radical, which then reacts further, Figure 112. The SH radical can also produce SH2 by reaction with itself, HO2, HCHO, H2O2 and CH3COOH. The OH reaction with COS is significantly slower than with CS2, so this is not a significant removal pathway in the troposphere13. The COS present is uniformly distributed, and has a long atmospheric lifetime (greater than 50 years), so it is also possible for some of this compound to reach the stratosphere. Unlike CS2, it is independent of the presence of O2 and of pressure13.
CH3S + O2 à CH3 + SO2 (6)
CH3 + NO2 à CH3SO + NO (7)
CH3SO + O3 à CH3SO2 + O2 (8)
CH3SO + O2 à CH3SO(O2) à products (mainly SO2) (9)
CH3SO + NO2 à CH3SO2 + NO (10)
OH + CH3SSCH3 à CH3SOH + CH3S (3)
OH + CH3SH à CH3S + H2O (4)
NO3 + CH3SH à CH3S + HNO3 (5)
OH + CH3SCH3 à H2O + CH2SCH3 (1)
OH + CH3SCH3 à CH3S(OH)CH3 (2)
DMS reacts initially with OH in two parallel steps (reactions 1-2)12. In air, the products react further with O2 to give either CH3SO3H (which dissociates to give CH3SO and HO2) or the CH3S radical, which then react further (reactions 6-10)12. DMDS is the dominant sulphur compound released from oceans. Both DMDS and methyl mercaptan, after initial reactions (3-5)13, produce products that go on to react as with DMS (reactions 6-10). These initial reactions all go via an intermediate complex. The DMDS reaction with OH is almost two orders of magnitude faster than that of DMS. This means that even small concentrations of this are significant, due to the higher reactivity13.
H2CO + HSO3- ßà CH2OHSO3- (17)
H2CO + SO32- ßà CH2O-SO3- (18)
SO2(g) + H2O ßà SO2.H2O(aq) (14)
SO2.H2O(aq) ßà HSO3- + H+ (15)
HSO3- ßà SO3- + H+ (16)
OH + SO2 + M à HOSO2 + M (11)
HOSO2 + O2 + M à HO2 + SO3 (12)
SO3 + H2O à H2SO4 (13)
Sulphur dioxide has mixing ratios of less than 10-9 in the free troposphere in remote areas; up to 30-10-9 in rural areas; and 2-10-6 in heavily polluted environments. This suggests an atmospheric lifetime of weeks12. In the troposphere, almost all of the SO2 is eventually oxidised to H2SO4. The H2SO4 forms aerosol particles which are rapidly incorporated into water droplets which then leave the atmosphere as acid rain. The threshold wavelength for the photodissociation is ca. 210nm. As these wavelengths do not penetrate the troposphere, it plays no part in the oxidation of SO212.
The only significant oxidant for SO2 in gaseous phase is the hydroxyl radical (reactions 11-13). HOSO2 can also react further with the hydroxyl radical to produce sulphuric acid. A minor reaction pathway of SO2 forms a complex with ammonia (sulfamic acid H2NSO3H, which then forms dimers). Although oxidation of SO2 by O3 alone is negligible, in the presence of alkenes, it is relatively fast. Water is an inhibitor in this system.
SO2 can form an aqueous equilibrium as well (reactions 14-16)14. The bisulfite ions produced can exist as either HOSO2- or HSO3-. Both these and sulphite ions can react with aldehydes when in solution (reactions 17-18). The product of reaction 17 is known as HMSA (hydroxymethanesulfonic acid), which is a stable and oxidant resistant species8. Oxidation of dissolved SO2 can also take place by O2 (catalysed by Fe3+/Mn2+), by O3, nitrogen oxides and by radicals such as OH and Cl. Likewise, it can also be oxidised in air on soot or graphite particles, as well as metal oxide surfaces. Water vapour enhances sulphate formation on surfaces13.
In conclusion, there are many different sulphur containing species in the troposphere, mainly emitted in the form of SO2, or oxidised to this form from either organic or inorganic compounds. All of the sulphur species are trace gases in the troposphere, but they have a profound impact on the chemistry here. The main removal mechanism is via sulphuric acid as acid rain.