10 ENVIRONMENTAL TECHNOLOGIES - treccani.it · The term nitrogen oxides is generally used to describe a mixture of nitrogen monoxide (95% of the total), nitrogen dioxide and traces

10

ENVIRONMENTALTECHNOLOGIES

10.1.1 Introduction

In recent years, growing concern about environmentalissues has focused attention on the quality ofatmospheric air, both in the troposphere, in otherwords from the ground up to an altitude of about 10km, and in the stratosphere, due to the well-known andworrying phenomenon of the decrease in theconcentration of ozone and the consequent thinning ofthe layer protecting the Earth from ultravioletradiation. This interest has resulted in the creation ofspecific environmental regulations, which made theirfirst appearance in the middle of the Twentiethcentury.

In Europe, this legislation sets out limiting valuesand objectives to be reached by a certain date in theform of Directives (among the Directives currently inforce it is worth recalling 96/62/EC, 99/30/EC,2000/69/EC and 2002/3/EC). Their adoption by eachmember state means that these Directives become statelaw. Additionally, on an international level,commissions have been established to deal withspecific problems and to draft protocols (the mostimportant include the Geneva Protocol of 1979, theMontreal Protocol of 1987 and the Kyoto Protocol of1997, with their subsequent amendments) subject toratification by individual countries. An attempt hasthus been made to take account of the impossibility ofrestricting to a single state or continent problemswhich inevitably affect neighbouring states (so-calledtransboundary pollution) or which, in the case ofextremely complex phenomena such as climate changeor the depletion of stratospheric ozone, affect theentire planet on a global scale. The latter twophenomena have been and remain (especially the first)an object of study by the scientific community and ofinterest to the authorities charged with managing theenvironment. The depletion of stratospheric ozone will

be dealt with below, as part of a brief overview of thecomplex interactions between the atmospheric ‘vessel’and the compounds of natural and man-made originemitted into it, thus creating a complex system. Thissystem can only be understood by taking into accountboth the chemistry of the aforementioned compoundsand the meteo-climatic factors which can alter theirconcentrations, for example by preventing their mixingand consequently their dispersion (episodes of thermalinversion are in fact closely correlated with cases ofextreme pollution).

There are four basic processes underlying pollutionphenomena: emission, transformation, diffusion andtransportation, and deposition.

The first two processes essentially concern thosepollutants described as ‘primary’, in other words thoseemitted directly as such, which may then undergotransformations and give rise to ‘secondary’ pollutants(such as nitrogen dioxide and ozone). These twoclasses of pollutants share their subsequent ‘itinerary’,which depends on the dynamic behaviour of the lowerlayers of the atmosphere, as far as transportation anddiffusion mechanisms are concerned, and is heavilyconditioned by this and by other meteo-climaticparameters (such as relative humidity), as far as sinkprocesses through dry and wet deposition areconcerned.

This chapter focuses on the chemistry of thecompounds of greatest environmental interest(nitrogen oxides, sulphur dioxide, non-methaneorganic compounds, ozone). Specifically, the reactionscharacterizing these in both the gas phase(homogeneous reactions) and in the aqueous phase(heterogeneous reactions) will be examined: thepresence of water in the atmosphere, in different statesof aggregation, is important for general considerationson the behaviour of these species after their releaseinto the atmosphere.

915VOLUME III / NEW DEVELOPMENTS: ENERGY, TRANSPORT, SUSTAINABILITY

10.1

Atmospheric chemistry

An in-depth understanding of all the processesgoverning pollution phenomena is essential both forthe ability to predict how these will evolve in the shortand long term, and to identify and implementappropriate control strategies.

10.1.2 Atmospheric chemistry in the gas phase

The apparent stability of the atmosphere derives fromthe fact that it is in a stationary state; this situation isdue to the relative constancy of inputs and outputswhich on average balance each other out. Since mostsources and sinks of the gases in the atmosphere arethe result of chemical reactions, an understanding ofreaction rates in the gas phase is extremely important.However, the study of chemical kinetics concentratesnot only on reaction rates but also on determining theprecise reaction mechanism.

This paragraph provides a brief survey of thereactions of the main pollutants in the gas phase.Specifically, Fig. 1 shows the reactions induced by thechemical species most heavily involved in theformation of acidic compounds in the atmosphere.

Nitrogen oxidesThe term nitrogen oxides is generally used to

describe a mixture of nitrogen monoxide (95% of thetotal), nitrogen dioxide and traces of dinitrogen oxide

(or nitrous oxide), nitrogen trioxide, tetroxide andpentoxide. Nitrogen monoxide is a colourless andodourless primary pollutant which forms fromcombustion at about 1,200°C, whilst nitrogen dioxideis a reddish pollutant with a strong and pungent smell.Nitrogen dioxide is a secondary pollutant since it doesnot form directly from combustion (except in apercentage of 4-5%).

The sequence of reactions leading to the formationof nitrogen oxides is described below.

Combustion produces NO (primary pollutant)through the reaction:

[1] N2�O2�� 2NO

which in turn can produce NO2 (secondary pollutant)via the (non-photochemical) thermal oxidationreaction:

[2] 2NO �O2�� 2NO2

However, reaction [2] is of minor importance sinceit is too slow at the concentrations of NO normallypresent in the atmosphere. The reaction kinetics is ofthe second order with respect to NO and of the firstorder with respect to O2; therefore, the kineticequation is R�k [O2] [NO]2, where R is the reactionrate and k is the rate constant (equal to 2.0 �10�38 cm6

molecules�2 s�1).Nitrogen dioxide is not produced only as shown

above; its formation is also triggered by ultravioletradiation and tropospheric ozone.

916 ENCYCLOPAEDIA OF HYDROCARBONS

ENVIRONMENTAL TECHNOLOGIES

nitrate aerosol

wet depositiondry depositionemission

aerosolHNO3NO2NO

O3

O3

NH3hn

OH

sulphate aerosol


NH3H2SO4SO2

H2O2

OH

ammonium aerosol


NHO3NH3

H2SO4

Fig. 1. Some reactionswhich occur in the atmosphere inducedby NOx, SO2 and NH3(the species most involvedin the formation of acid compounds in the atmosphere and in their subsequentsalification).

The photolysis of NO2 produces ozone through thereactions:

[3] NO2�hv (l �430 nm)�� NO �O(3P)

[4] O(3P) �O2�� O3

(k �6.0 �10�34 (T/300) �2,3 cm6 molecules�1s�1)

where hv is the light radiation and l is the wave length.The reaction between ozone and nitrogen

monoxide then produces nitrogen dioxide:

[5] O3�NO�� NO2�O2

(k �1.8 �10�14 cm3 molecules�1s�1)

There is not therefore a net loss of ozone, since thetitration reaction causes it to be destroyed whilst theNO2 photolysis reaction simultaneously causesreemission; it can thus be deduced that the threespecies are in fact involved in a photostationaryequilibrium. The rate constant in decompositionprocesses by photolysis of a trace species present inthe atmosphere, whose dimensions are the inverse oftime (s�1), can be generically defined as:

[6] J ��s(l)f(l)j(l)dl

where s(l) is the absorption cross-section expressed incm2 molecules�1, which is characteristic of eachchemical species, f(l) is the quantum yield of thephotolysis reaction, and finally j(l) is the actinic flux(number of photons�cm�2�s�1, generally used insteadof I�i�I�r�I�s�I�d); in other words it is the product of thecomponents of the solar radiation incident (I�i ) directlyon the air masses, consisting of reflected radiation (I�r ),scattered radiation (I�s ), and finally that which reachesthe Earth’s surface directly (I�d).

Assuming the absence of organic compounds inthe atmosphere, the relationship between theconcentrations of O3, NO and NO2 at a generic time tin a mass of air is a constant given by the ratio of JNO2

(the rate constant of NO2 photolysis) to kNO (the rateconstant of reaction [5] between NO and ozone):

JNO2[NO2]

[7] [O3]�11112kNO[NO]

The constant JNO2 varies depending on the angle ofthe luminous solar radiation, and so the relationshipbetween the concentrations changes over the course ofa day. The average value of JNO2 is 0.75�10�2 s�1 or0.533 min�1.

The presence of RO2� and HO2� radicals, whichform mainly during the hot season, promotes reactionswith nitrogen monoxide, in other words:

[8] RO2��NO�� NO2�RO(if R �CH3CH2CH2,

k �7.6 �10�12 cm3 molecules�1s�1)

[9] HO2��NO�� HO �NO2

(k �8.3 �10�12 cm3 molecules�1s�1)

As a result, under these conditions, reaction [5] ishindered and O3, being unable to react with NO,accumulates in the lower levels of the atmosphere.

In the presence of NO2, high concentrations of RO2�and HO2� free radicals give rise to alkyl nitratehydroperoxides, PeroxyAcetyl Nitrates (PAN) andPeroxyBenzoyl Nitrates (PBN).

PAN is a gas which tends to accumulatepersistently in the coldest parts of the troposphere; itthen spreads to warmer areas, leading to the formationof free radicals and NO2 since the dissociationconstant of PAN depends strongly on temperature.World Health Organization (WHO) guidelines statethat the annual averages of NOx in European cities arearound 40 mg/m3 and range from 20 to 90 mg/m3 inindustrialized countries; the base level is between 1and 9 mg/m3.

Nitrogen dioxide is more toxic than nitrogenmonoxide. For this reason, NO2 rather than NO ismonitored by law; values of 13 ppm cause irritation tothe mucous membranes of the nose and eyes; exposureto concentrations of 560 mg/m3 for 30 minutes causespulmonary problems. The WHO recommends thathourly mean concentrations of 200 mg/m3 and anaverage annual limit value of 40 mg/m3 should not beexceeded.

As already seen, NO2 is a pollutant which ismainly generated indirectly from the NO emitted bythe combustion of fuels used in road transportationand is thus described as a ‘mobile source pollutant’. Itis found mainly in urban areas with a high trafficdensity, coinciding with the opening and closing timesof workplaces and schools, and is also present in largecar parks. Another source of this substance iscombustion processes in civil and industrial plants.

As well as being harmful to human health, nitrogendioxide acidifies rain, degrades man-made objects,corrodes metals and damages vegetation.

During the day, NO2 is oxidized to nitric acid byreaction with the �OH radical:

[10] NO2��OH �M�� HNO3�M(k �1.1�10�11 cm3 molecules�1s�1)

where M is a third ‘body’ or molecule (typically N2 orO2) that has the role of removing the excess energythrough collision and so influences the kinetics of thereaction itself.

This reaction is slow compared to the NO-NO2

exchange: the average life-time of NOx is typically 1-2days in the low troposphere at middle latitudes. Thenitric acid is then deposited on the Earth’s surfacethrough various mechanisms.


ATMOSPHERIC CHEMISTRY

The NO3 radical is particularly important in thechemistry of organic compounds during the night,when its average life-time increases since it does notundergo photodissociation (as occurs during the day)and is thus able to react with Volatile OrganicCompounds (VOCs), oxidizing them rapidly. Its actionis similar to that of �OH during the daytime.

The reactions which take place can be grouped asfollows:• formation reactions:

[11] NO �O3�� NO2�O2

(k �1.8 �10�14 cm3 molecules�1s�1)

[12] NO2�O3�� NO3�O2

(k �3.2 �10�17 cm3 molecules�1s�1)

[13] NO2��OH�� HNO3

(k �1.1�10�11 cm3 molecules�1s�1)

• photodissociation reactions

[14] NO3�hv (l �640 nm)�� NO2�O(3P)

[15] NO3�hv (585 �l �640 nm)�� NO �O2

• reactions with organic compounds

[16] NO3�RH�� HNO3�R

[17] NO3�RCHO�� HNO3�RCO

[18] NO3�C�C�� C(ONO2)�C�

Addition reactions to alkenes (especially biogenicisoprene and monoterpenes) are much faster thanthose which occur with alkanes. The reaction of thenitrate with olefins also forms peroxide radicals (HO2��RO2�) and it has been shown (Salisbury et al., 2001)that NO3 may be as important in their formation as O3.

The formation of nitric acid due to reaction withhydrocarbons accounts for about 15% of the nitric acidin the atmosphere. Considering the averageconcentrations of these organic compounds in theatmosphere (in typically polluted urban atmospheresalkanes are present in concentrations of about 100 ppbwhilst formaldehyde and acetaldehyde are presentrespectively in concentrations of 20 ppb and 10 ppb)and the rate constants, and assuming that theconcentration of NO3 is 100 ppt, the net overall rate offormation is about 0.3 ppb h�1. Making similarassumptions for an average concentration of NO2 of50 ppb and �OH concentrations of 1�106 cm�3 (typicalof a moderately polluted atmosphere) the rate offormation is 2 ppb h�1.

The nitrate is thus a sink for NOx during the night,and may in turn be removed by various mechanismsinvolving other sinks such as reactions with organiccompounds or deposition on aerosols or on the ground(direct sink). It has in fact been shown that the averagelife-time of NO3 depends on the relative humiditydecreasing rapidly to less than 10 minutes when

relative humidity is 50% (Platt et al., 1984). This isprobably due to interaction with the water found on thesurfaces of the particles present in the environment.

In part, this may also be explained by the reactionof nitrogen pentoxide (N2O5) which represents anindirect sink for NO3 and which may react in thepresence of water, shifting the equilibrium:

[19] NO2�NO3�M��N2O5�M

For this reaction, the equilibrium constant reported inthe literature varies by a factor of 1.9 at ambienttemperature (from 1.8�10�11 to 3.44�10�11 cm3

molecules�1s�1).The N2O5 thus obtained reacts with water to give

nitric acid:

[20] N2O5�H2O�� 2 HNO3

(k �1.3 �10�21 cm3 molecules�1s�1)

Homogeneous hydrolysis thus represents a sink forNO3 and at the same time a source of nitric acid:although the rate constant seems fairly low, this reactioncontributes significantly to the formation of nitric acidin the homogeneous phase (0.3 ppb h�1 at 50% ofrelative humidity) as well as in the heterogeneousphase, and in this case is significantly faster.

Essentially, the main formation paths for nitric acidare the reaction between NO2 and �OH, the reactionbetween nitrate and organic compounds and thehydrolysis of N2O5. Nitric acid has a fairly longlife-time, and may thus be the terminal of variouschain reactions in the troposphere. It absorbs weaklyin the actinic region and thus does not undergophotolysis, however it may undergo deposition (dryand wet) and react with the �OH radical and (albeitslowly) with ammonia:

[21] HNO3��OH�� H2O �NO3

[22] HNO3�NH3�� NH4NO3

Ammonium nitrate is thus in equilibrium with thetwo species in the gas phase, and this equilibriumexists for both ammonium in the solid phase (relative humidity�62%) and in solution (relativehumidity�62%).

Sulphur dioxide (SO2)Sulphur dioxide (SO2) is a colourless gas with an

acrid, pungent smell. Emissions of this gas are mainlydue to the use of solid and liquid fuels and are directlycorrelated with their sulphur content:

[23] S �O2�� SO2

This is, therefore, a typical pollutant of industrialand urban areas, in the latter particularly during thewinter (due to domestic heating). Naturalconcentrations of SO2 are less than 5 mg/m3; annual



averages are lower than 50 mg/m3; and daily averagesrarely exceed 125 mg/m3.

Sulphur dioxide is used as a ‘tracer’, in otherwords a global indicator of atmospheric pollution, dueto its chemical stability in the atmosphere.

Most SO2 undergoes chemical transformationsbefore it reaches the ground; it is oxidized to SO3,

followed by hydrolysis to H2SO4, particles of whichabsorb further SO2, NH3 and traces of metals to form aparticulate aerosol which, depending on weatherconditions, may be transported for hundreds ofkilometres and reach the ground in the form of acid rain.

The oxidation reaction of SO2:

[24] 2SO2�O2�� 2SO3

has such a low rate in the absence of catalysts that itcan be completely disregarded as a source of SO3; thesame can be said for photooxidation as a reactionmechanism, since if every SO2

* molecule in an excitedstate was oxidized by reaction with O2 or other species,the average life-time of SO2 in the low tropospherewould be 52 minutes, which is, in fact, not the case.

In any case, the only rapid gas phase process whichis efficient enough to account for most of the sulphuricacid present in the aerosols formed by gas phaseprocesses is the reaction of SO2 with the �OH radical:

[25] SO2��OH �M�� HOSO2��M(k �1.1�10�12 cm3 molecules�1s�1)

It is known that a significant fraction of theHOSO2� radical is eventually transformed intosulphuric acid, but the rate of the reaction and itsintermediate products are not well known; as such, thereaction is generically described as:

[26] HOSO2�� H2SO4

In any case, a mechanism by which �OH isregenerated has been suggested:

[27] HOSO2�O2�� HO2��SO3

[28] HO2��NO�� NO2�OH

This mechanism was proposed by Calvert andStockwell (1984), on the basis of the experimentalevidence that in a photooxidant mixture of HNO2, NO,NO2 and CO, even the addition of substantial amountsof SO2 does not influence the concentration of �OH.

The oxidation of SO2 by this mechanism averagedout over 24 hours is 16.4%; during the winter the rateis lower due to the lower concentration of �OH.

Other oxidation reactions induced by other oxidizingspecies, such as O (3P), HO2� and CH3O2� , arecharacterized by lower rate constants than the reactionwith �OH, which remains the main oxidation reactioninduced by SO2 (the rate constants are k�5.7�10�14,k�1�10�18, and k�1�10�18 respectively).

Levels of SO2 are generally far higher than those ofSO3 since the latter, in contact with water vapour,leads to the following reaction, also encouraged by thepresence of particulate matter and solar radiation:

[29] SO3�H2O�� H2SO4

(k �9.1�10�13 cm3 molecules�1s�1)

Naturally, other fundamentally important oxidationreactions of SO2 to sulphuric acid are those which takeplace in the aqueous phase inside the water dropletspresent in the atmosphere; these will be dealt with inthe paragraph on solution equilibria.

Non-methane volatile organic compounds Methane is the most abundant hydrocarbon present

in the Earth’s atmosphere and the most stable withrespect to attack by �OH. This means that it can betransported far from its source before it is destroyed.By marked contrast, terpenes are extremely reactiveand as a result have short life-times. A broad variety ofhydrocarbons, the so-called Biogenic Non-MethaneHydrocarbons (BNHC), are emitted by naturalsources: these are unsaturated organic compoundsemitted mainly by plants, such as isopreneand some monoterpenes like a-pinene, b-pinene,d-limonene, etc.

In towns the air may contain variableconcentrations (even above 1.0-2.0 mg/m3) ofhydrocarbons other than methane, whilst methaneitself may exceed 1.5 mg/m3.

The most important reactions involving thesespecies (RH) are those which lead to the formation ofozone:

[30] RH��OH �O2�� RO2�(�H2O)

[31] RO2��NO�� RO �NO2

[32] NO2�O2�hv�� O3�NO

As such, the amount of ozone produced per moleof hydrocarbon depends on its concentration and itsreactivity with �OH.

The relative reactivity of species j in a mixture of ihydrocarbons (HC) can thus be expressed as:

kj[HC]j[33] Rj �

111133

�i

ki[HC]i

Taking account of the typical abundances of thesecompounds and their reactivity, it is possibleto calculate the rates of degradation. For example,Table 1 shows the removal rates of some of the mostabundant hydrocarbons in the atmosphere.

In terms of the rate of degradation, alkenes havethe highest potential for forming ozone, followed byaromatic compounds, whilst alkanes have the lowestformation potential.



The oxidation of hydrocarbons is strictly correlatedwith the photostationary state NOx�O3 described bythe Leighton relationship: the oxidation of VOCs(Fig. 2) by �OH produces the radical RO2� whichconverts NO to NO2 through the reaction alreadydescribed for NOx without involving the simultaneousremoval of O3. As a result, the concentration of ozoneincreases, forming other �OH radicals, thus increasingthe rate at which VOCs are oxidized.

The maximum production of O3 requires anappropriate ratio of NOx to VOC concentrations: a lackof NOx results in an insufficient production of �OH toinduce the oxidation of VOCs, whilst a lack of VOCsmakes it impossible to reach the concentrations ofRO2� needed to alter the photostationary statesignificantly.

The non-linearity of ozone formation has beendemonstrated experimentally by Kelly and Gunst (1990).

Surveying the main reactions of the various classesof organic compounds, it can be seen that the main

reactions are those with �OH and NO3, whoseconstants for the various hydrocarbons are reported inTable 2.

The chemistry of the higher term alkanes followsthe same mechanism as methane: the attack by �OHtakes place preferentially to form the most stable alkylradical, therefore tertiary and secondary hydrogens arethose which react most easily with �OH. The chemistryof anthropogenic alkenes follows a similar mechanismto that of biogenic isoprene: the initial reaction is theaddition of �OH followed by the addition of O2 to ahydroxy-substituted alkylperoxy radical, which in turnreacts with either NO or HO2� depending on whetherthe concentrations of NO are high or low.

The photochemical oxidation of carbonylcompounds leads to the production of peroxyacetylradicals and the following reactions give rise to theformation of PAN (CH3C(O)O2NO2):

[34] CH3CHO ��OH (�O2)�� CH3C(O)O2�H2O(k �1.6 �10�11 cm3 molecules�1s�1)

[35] CH3C(O)O2�NO2�M�� CH3C(O)O2NO2�M(k �3.6 �10�12 cm3 molecules�1s�1)



VOC

RO2�

�OH, O2, M

O2

O2

NO

hn

O2

NO

via aldehydes

RO�NO2

O

O3

�RCH3

�RCH2O2 R�CHO

� RCHO�RCH2O

hn

Fig. 2. Schematic of the oxidation of organic compounds in the troposphere.

Table 1. Kinetic data on the removal of hydrocarbons present in the atmosphere by �OH radicals

COMPOUNDk at 298K

(10�12 cm3, molec.�1s�1)Concentration

(1010 molec. cm�3)Removal rate

(10�2s�1)

Methane 0.0077 5,748.0 44.3

Toluene 6.4 12.1 77.4

Ethylene 8.8 26.8 235.8

Acetylene 0.9 16.1 14.5

Benzene 1.0 5.3 5.3

Table 2. Rate constantsfor VOC oxidation reactions

COMPOUNDCONCENTRATION

(ppb carbon)kOH(1012cm3

molec.�1s�1)kNO3

(1016cm3

molec.�1s�1)

Isopentane 45.3 3.9 1.6

Toluene 33.8 5.96 0.3

Ethylene 21.4 8.52 2.1

Acetylene 12.9 0.9 �0.2

Benzene 12.6 1.23 0.2

Isoprene – 101 5,900

a-pinene – 53.7 58,000

[36] CH3C(O)O2NO2�M�� CH3C(O)O2�NO2�M(k �1.8 �10�4 cm3 molecules�1s�1)

The formation of PAN is a process which ends thepropagation of the reaction chain, and which competeswith the reaction between the peroxy radical and NO:

[37] CH3C(O)O2�NO�� CH3C(O)O �NO2

(k �1�10�11cm3 molecules�1s�1)

[38] CH3C(O)O �O2�� CH3O2��CO2

(k �2.0 �10�12 cm3 molecules�1s�1)

[39] CH3O2��NO�� CH3O �NO2

(k �7.6 �10�12 cm3 molecules�1s�1)

[40] CH3O �O2�� HCHO �HO2�(k �1.9 �10�15 cm3 molecules�1s�1)

The formation of PAN is encouraged by lowtemperatures and pressures. Thermal decomposition isthe most important destruction path for PAN near theEarth’s surface, whilst at altitudes above 7 km it reactswith �OH:

[41] CH3C(O)O2NO2��OH�� products(k �1.1�10�13cm3 molecules�1s�1)

Therefore, if the PAN formed rises rapidly toheight atmosphere, its life-time increases and it mayrepresent a source of NOx through long-distancetransportation mechanisms.

Ozone and photochemical smogFor a photochemical smog process to be triggered,

sunlight, nitrogen oxides and volatile organiccompounds must be present; additionally, the processis favoured by a high atmospheric temperature. Sincenitrogen oxides and volatile organic compounds areamong the main components of emissions in urbanareas, towns located in geographical areascharacterized by intense solar radiation and hightemperatures (such as those in the Mediterranean) areideal candidates for episodes of intense photochemicalpollution. The knowledge necessary for understandingsecondary pollution events thus concerns the chemicaland chemico-physical transformation processesundergone by pollutants, dynamic processes in thelower atmosphere (atmospheric stability, direction andintensity of the wind) and the intensity of solar radiation.

Ozone is a photochemical oxidant similar to PAN,nitrogen dioxide and hydrogen peroxide which areessentially secondary pollutants formed in thetroposphere by chemical reactions starting fromprimary pollutants (basically VOCs and nitrogenoxides which are therefore also described asprecursors) in the presence of solar radiation. It isnaturally present in the troposphere at concentrationsranging from 20 to 80 ppb.

Ozone has a characteristic smell and may causesevere irritation to the respiratory system and the eyesat concentrations exceeding 200 ppb. It is also a causeof the oxidative degradation of some non-biologicalmaterials, especially elastomers, textile fibres anddyes.

In fact, the presence of only six electrons in thevalence shell of oxygen gives it electrophilicproperties and therefore the tendency to removeelectrons from other species or to share them. It ischaracterized by a redox potential of 2.07 V in anaqueous system.

The formation reactions of ozone by nitrogenoxides and hydrocarbons have already been reportedextensively in the chapter devoted to those pollutants.

In the lower atmosphere, ozone forms from thereaction of atmospheric oxygen with the atomicoxygen produced by the photolysis of nitrogendioxide; the ozone formed is in turn removed bynitrogen monoxide, with the new formation of NO2.

In unpolluted atmospheres, where other chemicalspecies are not present in appreciable quantities, thisseries of reactions forms a cycle (the photostationaryozone cycle) and there is no possibility forphotochemical pollution. The fundamental steptowards the atmospheric enrichment of ozone andother photooxidant species (in other words oxidizingchemical species formed by chemical reactions thatoccur only in the presence of light) is the formation ofNO2 by alternative pathways which do not entail theremoval of ozone. Identifying the formation pathwaysof NO2 thus represents the key to understandingphotochemical oxidation processes.

The main alternative pathway for the formation ofNO2 is the oxidation of NO by peroxide radicals(RO2�). These free radicals originate from thedegradation of volatile hydrocarbon molecules (RH)and their subsequent reaction with atmosphericoxygen. The attack on volatile hydrocarbons is due tothe presence in the atmosphere of other free radicals,�OH hydroxyl radicals: the processes which generatehydroxyl radicals are thus fundamental in triggeringphotochemical pollution processes. The production of�OH radicals is essentially a photochemical processand the main precursors are nitrous acid, formaldehydeand ozone itself. Ozone is therefore not only the mostquantitatively important product of photochemicalpollution processes, but also part of the ‘fuel’ whichactivates the process.

In fact, ozone, like nitrogen dioxide, undergoesphotolysis and given the energy of the O2�O bondwhich is only 101 kJ mol�1 (in other words in theorder of 1 eV per molecule) and since the energy E ofa photon is linked to frequency n by the relationshipE�hn, fairly low frequency radiation is needed (but in



any case over 2.53�1014 Hz, in other words with awavelength shorter than 1.185 nm) to cleave the bond;most of the spectrum of solar radiation thus hassufficient energy to cleave the O2�O bond:

[42] O3(1A1) �hv (l�1,180 nm)�� O(3P) �O2(3Sg�)(J �4.2 �10�4 s�1)

At shorter wavelengths, the photon’s excess energymay be converted into the electronic excitation of theproducts:

[43] O3(1A1)�hv (l�310 nm)�� O(1D) �O2(1Dg)(J �2.9 �10�5 s�1)

Electronically excited molecular oxygen (1Dg) maybe another potential oxidizing agent in thetroposphere, especially for unsaturated hydrocarbons:the rate constants for reactions with these compoundsare in the order of 10�18 cm3 molecules�1s�1.Comparing this with the rate constants of the reactionsof alkenes with �OH and O3 (on the order of 10�12, asreported in Table 1, and 10�18 cm3 molecules�1s�1

respectively) leads to the conclusion that, to have asignificant effect on tropospheric chemistry, theconcentrations of O2 (1Dg) would need to be of thesame order of magnitude as ozone or greater. In fact,concentrations of this species are around 10 ppt.

Since the transition:

[44] O (1D)�� O (3P)

is impossible, as this is a spin-forbidden transition, O(1D) may undergo thermal degradation throughcollisional energy transfer or react with other species,like CH4 or H2O, extracting a proton from them andthus giving rise to �OH radicals:

[45] O(1D) �H2O�� 2�OH(k �2.2 �10�10 cm3 molecules�1s�1)

This reaction is extremely fast and occurs incompetition with the deactivation of the O (1D) by air(indicated by M):

[46] O(1D) �M�� O (3P) �M(k �2.9 �10�11 cm3 molecules�1s�1)

In an atmosphere characterized by a relativehumidity of 50% and a temperature of 298K, about10% of the O (1D) produced reacts with water to formhydroxyl radicals.

At middle latitudes, the main source of �OHradicals is ozone, given the high concentrations of thisspecies which is about 40 ppb or 9.84�1011 moleculescm�3.

In fact, the rate of production of �OH radicals fromthe photolysis of ozone at middle latitudes is:

[47] R(�OH, O3) �4.65 �105 molecules cm�3 s�1

Other sources of �OH radicals are the photolysis ofnitrous acid, hydrogen peroxide and formaldehyde.

Nitrous acid and formaldehyde are precursors of�OH radicals, but in their turn have formationpathways which are essentially secondary, startingfrom species involved in photochemical processes(nitrogen dioxide for nitrous acid and hydrocarbonsand radicals or ozone for formaldehyde).

The formation reactions for �OH radicals startingfrom nitrous acid and hydrogen peroxide are as follows:

[48] HONO �hv(l �320 nm)�� OH �NO(J �1.8 �10�3 s�1)

[49] H2O2�hv(l �360 nm)�� 2�OH(J �6.9 �10�6 s�1)

In the case of formaldehyde, the predominantmechanism when 300�l �320 nm is:

[50] HCHO �hv�� H �CHO (J �1.7 �10�5s�1)

Subsequent oxidation leads to the formation of the�OH radical:

[51] H �O2�M�� HO2��M(k �7.0 �10�13 cm3 molecules�1s�1)

[52] CHO �O2�� HO2��CO(k �6.0 �10�12 cm3 molecules�1s�1)

[53] HO2��NO�� OH �NO2

(k �8.0 �10�12 cm3 molecules�1s�1)

At wavelengths above 340 nm there is apreferential dissociation into relatively stable products:

[54] HCHO �hv�� H2�CO (J �4.3 �10�5s�1)

Crutzen (1988) has estimated that on average50-60% of the formaldehyde follows the mechanismabove, 20-25% follows the mechanism leading to theformation of CHO and H and subsequent oxidationreactions, and 20-30% reacts directly with �OHaccording to the reaction:

[55] HCHO ��OH�� HCO �H2O(k �1.0 �10�11 cm3 molecules�1s�1)

The �OH radical may also give rise to the oxidationof CO:

[56] CO ��OH�� CO2�H(k �2.0 �10�13 cm3 molecules�1s�1)

[57] H �O2�M�� HO2��M(k �7.0 �10�13 cm3 molecules�1s�1)

In the presence of a source of NO, �OH willtherefore be obtained again directly via the oxidationreaction induced by the HO2� radical, and indirectly viathe photolysis of the ozone produced.

The reactions involving formaldehyde also producehydroperoxide radicals, which are among the radical



species fundamental in photochemical smog processesalongside hydroxyl radicals and alkyl peroxideradicals (formed, as seen above, from organiccompounds).

There is an interconversion between hydroxylradicals and hydroperoxide radicals, and the keyreaction in this process is as follows:

[58] HO2��NO�� OH �NO2

Photochemical pollution is thus caused by asequence of interdependent reactions (in some casestrue chain reactions) which give rise to a processwhich feeds itself; this explains why acute episodes ofphotochemical smog often last for several consecutivedays, increasing in intensity.

Atmospheric secondary particulate matter The formation of aerosols in the atmosphere has a

significant impact on visibility, the climate and thechemical processes which occur in the atmosphere; itis also of special interest since the finest fraction (withan aerodynamic diameter less than 2.5 mm) maypenetrate the alveoli of the lungs and thus has a directimpact on human health.

Aerosols in the troposphere may be emitteddirectly (primary particulate matter) or be formedfrom chemical processes (secondary particulatematter). The sources are both natural and man made,and the composition of this material therefore varies

considerably. The dimensions of the particles varysignificantly for both primary and secondaryparticulate matter, with aerodynamic diametersranging from 2 nm to over 10 mm.

Secondary aerosols are generated by the gas-particle conversion which follows the formation,through oxidative processes, of products characterizedby particularly low volatility or high solubility. Sincethese oxidative processes are often of a photochemicalnature, the resulting aerosols may be counted amongthe secondary photochemical pollutants.

Secondary aerosol may be generated either bycondensation onto existing aerosol or by nucleation toform new particles or suspended droplets (Seinfeld,1986; Clement and Ford, 1996). The most importantmolecule which may give rise to the nucleation of newparticles is H2SO4.

Fig. 3 summarizes the processes which occur inthe presence of primary and secondary pollutantsand UV radiation: the interactions between the latterand the particles have been the object of recentstudies (Kikas et al., 2001). These have shown thatthe rate of photolysis of some pollutants (O3, NO2)depends on the presence of particles which absorbUV radiation (especially elementary carbon andorganic aerosols; Jacobson, 1999) or scatter it (theoptical properties of the particles vary with size andare strictly linked to their chemical composition). Astudy by Jacobson (1998) has shown that levels of



UV radiation

(NO3, HONO, PAN, RO2�, etc.)

scattering gases

absorbing gases

scattering particulate matter

absorbing particulate matter

HONO VOCNO2O3

semi-volatile organics

H2SO4

particulatesulfate

secondaryorganic aerosol

particulatenitrate

SO2

NH3

HNO3

O3, NO3, HO�

HO�, aqueousoxidation

HO�

Fig. 3. Negative feedbackcycle between UV radiationand particulate matter.

ozone in Los Angeles have fallen by 5-8% due toaerosols. Wendisch and colleagues (1996) havemeasured the vertical profile of particulate matter tocalculate the vertical rate profile of the NO2

photolysis reaction.With regard to the composition of secondary

particulate matter, studies have been undertakenexploiting the different optical properties (lightscattering) at different relative humidities of thespecies of interest: H2SO4 and (NH4)2SO4.

These studies have shown that ammonium sulphateis often observed in real samples, and there istherefore sufficient ammonia in ambient air tocompletely neutralize sulphuric acid (Weiss et al.,1977). However, sulphate may also be present in otherforms, including ammonium disulphate(NH4)3H(SO4)2; the formation of mixed salts withnitrate is unproven, but seems probable under normalatmospheric conditions.

The equilibria characterizing mixtures of solidsand solutions containing NH4NO3 or mixtures of NH4

�,NO3

� and SO42� in equilibrium with NH3 and HNO3 at

the concentrations typical of the atmosphere have beenstudied and are reported below:

[59] NH3(g)�H2SO4(g)��

��NH4��HSO4

�

[60] 2NH3(g)�H2SO4(g)��

��2 NH4��SO4

2�

[61] NH3(g)�HNO3(g)��

��NH4��NO3

�

[62] NH3(g)�H2SO4(g)��

��NH4HSO4(s)

[63] 2NH3(g)�H2SO4(g)��

��(NH4)2SO4(s)

[64] NH3(g)�HNO3(g)��

��NH4NO3(s)

[65] 4NH3(g)�2 HNO3(g)�

�H2SO4(g)��

��(NH4)2SO4�2NH4NO3(s)

[66] 5NH3(g)�3 HNO3(g)�

�H2SO4(g)��

��(NH4)2SO4�3NH4NO3(s)

Examining the composition of particulate matterindicates the prevalence of NH4

� and NO3� in the

largest size fraction (diameter between 0.1 and 1 mm);this is due to the Kelvin effect which involves thecreation of a greater vapour pressure of the volatilespecies NH3 and HNO3 on strongly curved surfaces.This leads to the volatilization of NH3 and HNO3 fromthe smallest particles and condensation onto thelargest; this phenomenon does not occur for sulphate,which has a low volatility.

Fig. 4 shows the distribution of the various speciespresent in particulate matter as a function of theequivalent diameter of the particles themselves. Othercomponents of particulate matter are metals andorganic compounds.

Most Organic Carbon (OC) is found in the finefraction of particulate matter. OC derives mainly fromthe oxidation of combustion products, such as VOCs,and their subsequent condensation, dissolution into theaqueous phase, adsorption (especially onto particles ofelementary carbon, EC), or absorption (Seigneur,2001). The OC found in the particulate matter emittedby motor vehicles contains more than 100 differentcompounds, including alkanes, benzaldehydes, andPolycyclic Aromatic Hydrocarbons (PAHs),particularly dangerous to human health (Rogge et al.,1993).

The organic compounds of greatest medicalinterest are the polycyclic aromatic hydrocarbons:these form from hydrocarbons with a low molecularmass by pyrosynthesis at temperatures over 500°C.The result is the formation of several aromatic ringscondensed into very stable structures which in theatmosphere, as an effect of solar radiation, may betransformed into more dangerous compounds such asnitro-PAHs by reaction with nitric acid and oxidizedPAHs by reaction with ozone. The most frequentlymentioned PAH is benzo(a)pyrene, which



NH�

dC (

equi

vale

nt /

m3 )

/ dD

ae (

mm

)

500

0

100

200

300

400

equivalent diameter Dae (mm)10�2 10�1 101

NH42�SO4

NO�NO3

Na�

Cl�

H�

Fig. 4. Distribution of concentration of ions with respect to equivalent diameter.

becomes highly carcinogenic through metabolicactivation. High concentrations of PAHs are present inthe soot generated by the combustion of biomassesand coal, and by the exhausts of diesel and gasolinevehicles.

10.1.3 Atmospheric chemistry in the aqueous phase

BasicsThe total volume of water in the atmosphere is

estimated to be about 1.3�1013 m3; this water is presentin different forms: aerosols, clouds, fog and rain.

At the northernmost latitudes, over 30% of thelower troposphere is occupied by cloud bodies. Thewater content of a cloud is in the order of 0.1-1 g m�3

and the size of the droplets forming it depends on thetype of cloud; in general, their diameter is greater than10 mm. Fog has a lower water content, about 0.1 gm�3, and smaller droplets.

Aqueous aerosols in the atmosphere consist ofparticles with a broad range of diameters, on the basisof which it is classified as a fine or coarse fraction. Afine fraction contains free water in various forms:dilute aqueous solutions, supersaturated solutions andfine films on insoluble particles.

The volume of the drops present in fog and cloudsis the liquid medium on which the absorption ofreactive trace gases occurs; the majority of these arehighly soluble in water. The concentrations of somereactive species may therefore be higher in water thanin the surrounding air. This, in conjunction with thehigh reaction rates of some species in the aqueousphase, leads to the conclusion that the dropletscontained in fog and clouds may be extremely efficientreactors for the oxidation of SO2, NO and NO2.

An important part of atmospheric chemistry takesplace on suspended particles or droplets. The reactionswhich occur on the surface or inside these particles aredescribed as heterogeneous: they take place on theinterface between two phases, i.e. gas-liquid andgas-solid. Those which take place internally ratherthan on the surface are not heterogeneous in the strictsense of the word, but can be regarded as suchconsidering the volume of air containing the particles.

The water content in the atmosphere can beexpressed in terms of grams (or of cm3) of water perm3 of air or as the adimensional fraction of volume L(for example, m3 of water per m3 of air). Values of L indifferent forms of water condensation in theatmosphere are:• Clouds, L�10�7-10�6

• Fog, L�5�10�8-5�10�7

• Aerosols, L�10�1-10�10

The penetration of atmospheric gases into thesuspended droplets involves the following steps:• Transportation of the gas species towards the

surface of the drop. • Absorption and transportation across the air-drop

interface. • Transportation into the body of the drop. • Reactions inside the drop.

The solubility of gases in water is described byHenry’s law, which states that at equilibrium the partialpressure of a gas on a solution containing the same gasis proportional to its concentration in solution.

The absorption of a generic species, A, into watercan be represented by the following, entirelyequivalent, equations:

[67] A(g)�H2O��A�H2O

[68] A(g)��

��A(aq)

The equilibrium between A in the gaseous formand the same species in solution can be expressed byHenry’s constant KH

[69] KH�[A(aq)]�pA

where pA is the partial pressure of A in the gas phaseand [A(aq)] is the concentration of A in the aqueousphase at equilibrium. The measurement units forHenry’s constant are [mol l�1 bar�1]. Typical values ofHenry’s constant for the main atmospheric gases aregiven in Table 3. The higher the value of this constant,the more soluble the gas in the aqueous phase;however, it is important to remember that some gasessubsequently react with the water. Henry’s constantonly considers solubility (physical process) and nothydrolysis reactions (chemical process).

The effect of equilibria in solution is to increasethe quantity of gas passing from the gas phase to theliquid phase in the atmosphere compared to levelspredicted based on Henry’s law alone.

Given the rapid development of acid-baseequilibria, Henry’s law is often extended by defining aHenry’s pseudoconstant KH* which takes into account



Table 3. Values of Henry’s constantfor the main atmospheric gases

GasKHat 298K

(mol l�1 bar�1)

Oxygen 1.3 �10�3

Ozone 9.4 �10-3

Nitrogen dioxide 1�10�2

Carbon dioxide 3.4 �10�2

Sulphur dioxide 1.24

Ammonia 62

all dissolved species. For example, in the case of SO2

we can define:

[S(IV)] [SO2(aq)]�[HSO3

�][70] KH*

S(IV)�11144�12111111244�pSO2

pSO2

K1 K1K2�KHSO2�1�11�112�[H�] [H�]2

where K1 and K2 are the first and second dissociationconstants of the sulphurous acid. It can be seen fromthis equation that the value of KH* depends on the pH:the solubility of SO2 decreases as the pH decreases.Despite its high solubility, sulphur dioxide is not foundcompletely dissolved in the tiny water droplets whichform clouds.

The distribution of a generic species, A, betweenthe gas and aqueous phases in a cloud can beexpressed in terms of the ratio of concentrations of Ain the two phases per unit volume of air:

moles of A in solution per litre of air[71] 1111121121111111244�

moles of A in air per litre of air

KH�pA �L�11112�KH�R�T �L

pA �R �T

The constant KH is indicated as KH* if the speciesparticipates in dissociation equilibria in solution.

If KH�R�T�L1 or KH1/RTL, species A is presentpredominantly in the gas phase, whilst the contrary istrue if KH1/RTL. Therefore, if L�10�6,1/RTL�4�10�4 mol l�1 bar�1: if Henry’s constant for aspecies is less than 4�10�4 mol l�1 bar�1 it will bepresent mainly in the gas phase.

If the pH is 4 and L is 10�6, the value ofKH

S(IV)�102 gives KH*S(IV)1/RTL, and SO2 is therefore

present mainly as a gas inside clouds. By contrast, fora species such as HNO3, KH*�1010 and therefore atequilibrium this species will be present almost entirelyin solution. In acidic environments, H2O2 and NH3 arealso found mainly in the aqueous phase.

The equilibria for the most important species inatmospheric chemistry in the aqueous phase can bedescribed as follows:• carbon dioxide

[72] CO2�H2O��H2CO3(aq)

[73] H2CO3(aq)��

��H�(aq)�HCO3

�(aq)

[74] HCO3�

(aq)��

��H�(aq)�CO2�

3(aq)

• sulfur dioxide

[75] SO2(g)�H2O��H2SO3(aq)

[76] H2SO3(aq)��

��H�(aq)�HSO�

3(aq)

[77] HSO�3(aq)�

��H�

(aq)�SO2�3(aq)

• ammonia

[78] NH3�H2O��NH4OH(aq)

[79] NH4OH(aq)��

��NH�4(aq)�OH�

(aq)

Henry’s constant KH and the equilibrium constantsfor reactions in the aqueous phase (K� and K�) areprovided in Table 4.

Both in the case of sulphur dioxide and carbondioxide, the second dissociation constant is muchsmaller than the first dissociation constant, and cantherefore be disregarded at low pH values.

The pH of a water droplet in equilibrium withatmospheric CO2 can be calculated by combining thetwo equilibrium constants referring to the solubilityand the first dissociation:

[80] [HCO�3 ] �KH�K��pCO2

�[H�]

If the only source of hydrogen ions is thedissociation of carbon dioxide, [HCO�

3 ] = [ H�] andtherefore:

[81] [HCO�3 ] �[H�] �(KH �K��pCO2

)1/2

Assuming a partial pressure of CO2 of 340 ppm,this gives a pH of 5.6. This is the pH of rainfall inremote areas; the presence of traces of othercompounds may affect acidity: SO2 present at a levelof 5�10�9 bar gives a pH for the solution of 4.6. Infact, for a solution in equilibrium with SO2, it ispossible to write an equation similar to that for carbondioxide:

[82] [HSO�3 ]�(KH �K��pSO2

)1/2

It therefore seems obvious that even lowconcentrations of SO2 have a profound effect on pHalthough the concentration of CO2 in the atmosphere isfar higher: this is due to sulphur dioxide’s higher



Table 4. Values of Henry’s constant and equilibrium constants in the aqueous phase for NH3, SO2 and CO2

Gas KHat 288K (mol l�1 bar�1) K�(mol l�1) K�(mol l�1)

Ammonia 90 1.6�10�5

Sulphur dioxide 5.4 2.7�10�2 1�10�7

Carbon dioxide 0.045 3.8�10�7 3.7�10�11

solubility and dissociation constant. These propertiesgive this species a greater acidifying power which mayeven be intensified by concomitant S(IV)�� S(VI)sulphur oxidizing reactions.

Sulphur oxidationThe oxidation of SO2 in the gas phase is mainly

induced by the �OH radical:

[83] �OH�SO2 (�M)�� HOSO2 (�M)

[84] HOSO2�� H2SO4

The �OH radical forms in the atmosphere in thepresence of mixtures containing non-methanehydrocarbons and NOx due to the effect of solarradiation and gives rise, through the formation of theaforementioned radical species, to the formation ofsulphuric acid. There are various hypotheses regardingthe mechanism by which the radical species formedgives rise to H2SO4 (Benson, 1978; Calvert et al.,1978; Davis et al., 1979); this discussion consequentlymakes use of the simplified equation above.

Oxidation in the gas phase induced by the �OHradical is characterized by formation rates greater than1% an hour; studies conducted both in the field and inthe laboratory have recorded values for the conversionrate higher than those predictable on the basis of gasphase chemistry alone. This leads to the conclusionthat reactions in the aqueous phase may be importantsources of sulphate and in some cases may even bemore important than those in the gas phase.

When SO2 dissolves in an aqueous solution, thethree resulting species, SO2�H2O, HSO3

� and SO32�,

with S(IV), may be oxidized by various species:oxygen (catalyzed by iron and manganese), ozone (thedominant mechanism when pH�5) and hydrogenperoxide (the dominant mechanism when pH�5).

Fig. 5 shows schematically the distributionequilibrium and the subsequent oxidation equilibria ofsulphur dioxide.

The main oxidation mechanisms are:• ozone

[85] HSO�3(aq)�OH��O3�� SO4

2��H2O �O2

• hydrogen peroxide

[86] HSO�3(aq)�H2O2(aq)�

��SO2OOH��H2O

[87] SO2OOH��H�� H2SO4

• oxygen

[88] HSO�3(aq)�1�2O2(aq)��

Fe, MnSO2�

4(aq)�H�(aq)

Ozone reacts very slowly with SO2 in the gasphase, whereas the reaction is fast in the aqueousphase. The most plausible mechanism is the ionicmechanism shown above and proposed by Maahs(1983). Radical mechanisms have also beenproposed whose contribution to the oxidation ofS(IV) is not quantifiable with any degree ofcertainty. These may take place via a radicalintermediate such as �OH (Hoignè and Bader, 1975;Penkett et al., 1979):



(2)

SO2(g)

SO2(interface)

SO2�H2O

transport across air-water interface

(4) transport into bulk phase

(5) oxidation

(3) establishment of S(IV) equilibria

S(IV)S(VI)

(1) transport to droplet surface

SO32�

HSO3�

Fig. 5. Sulphur dioxidedistribution and oxidationequilibria.

[89] HSO3��OH�� H2O �SO3

�

[90] SO3��O2�� SO5

�

[91] 2SO5�� SO4

��SO42��O2

Oxidation induced by H2O2 is relativelyindependent of pH and this is due to the fact that therate constant of the reaction and the solubility of S(IV)show opposing trends with respect to pH. Otherspecies are affected in different ways by variations inpH: in the case of ozone the reaction rate increases byone order of magnitude when passing from pH 1 to pH3, whereas the reaction rate of oxygen catalyzed byiron increases by two orders of magnitude. Thedependence of rate and the solubility of S(IV) on pHmeans that the reaction with ozone is only importantwhen pH�5.

Research has concentrated particularly on reactionsin the aqueous phase catalyzed by transition metals:the most important catalysts are iron (in the formFe3�) and manganese (in the form Mn2�). Somestudies have been carried out on the possible catalyticaction of other metals such as Cu2� and Co2�;however, the reactions catalyzed by the latter arecharacterized by very low reaction rates under typicalambient conditions.

The rate of the reaction catalyzed by Fe(III)depends on various parameters, such as pH, ionicforce and temperature, and is affected by the presenceof some anions (such as SO4

2�) and cations (Mn2�) insolution.

The mechanisms of these reactions have long beendebated. A radical mechanism has been suggested(Hoffmann and Boyce, 1983; Hoffmann and Jacob,1984):

• initiation

[92] Mn��SO32�� M(n�1)��SO3

��

• propagation

[93] SO3�� O2�� SO5

��

[94] SO5�� SO3

2�� SO52��SO3

��

• oxidation

[95] SO52��SO3

2�� 2SO42�

• termination

[96] 2SO3�� S2O6

2�

[97] SO3�� SO5

�� S2O62��O2

[98] SO5��SO5

�� S2O82��O2

This hypothesis was later abandoned since itclashed with the kinetic order of reaction of thereagents and was replaced by two alternativemechanisms: ionic and photochemical.

The ionic mechanism, initially proposed by Bassettand Parker (1951), involves first the formation of a

metal-sulphite complex followed by bonding withoxygen:

[99] Mn2��2SO32��

��Mn(SO3)22�

[100] Mn(SO3)22��O2�

��Mn(SO3)2O2

2�

[101] Mn(SO3)2O22��

��Mn2��2SO42�

Mechanisms that are even partially photochemicalhave been suggested, based on the observation (Lunakand Veprek-Siska, 1976) that the oxidation of S(IV) inhomogeneous aqueous systems for wavelengths above300 nm does not occur unless Fe3� is present.Photooxidation has been attributed to the absorption oflight by a Fe3�- S(IV) complex.

In the case of iron, this is present in solution in thesoluble and solid form:

[102] [Fe]tot�[Fe]sol�[Fe(OH)3]

Soluble iron is found in aqueous solution invarious ionic forms:

[103] [Fe]sol�[Fe3�] �[FeOH2�] �[Fe(OH)2�] �

�[Fe2(OH)24�]

The equilibrium between iron in the soluble formand solid iron is as follows:

[104] Fe(OH)3�3H��Fe3��3H2O

As a result, the relative quantities of the differentforms of iron in aqueous solution are stronglyinfluenced by pH. With pH values of 4.5,[Fe3�]�3�10�11; above this pH value the soluble formof iron decreases significantly.

Another catalyst for these oxidation reactions ismanganese: attempts have been made to determinewhich of the two metals is the main cause of SO2

oxidation, but to date their relative importance in thereaction has not been clarified.

It is likely that the two catalysts act in synergy: anincrease in the reaction rate has been observed in thepresence of both ions, higher than would be expectedon the basis of the sum of their respective catalysis



Table 5. Concentrations (expressed in molarity, M)of iron and manganese in various aqueous matrices

present in the atmosphere

AQUEOUS MATRIX Mn Fe

Mist 10�7-10�4 10�4-10�3

Clouds 10�8-10�5 10�7-10�4

Rain 10�8-10�6 10�8-10�5

Fog 10�7-10�5 10�6-10�4

rates: the removal of S(IV) is about 3–10 times fasterif there is synergy between the two species (Martin,1984). Table 5 quantifies the presence of these twospecies in the various forms of water condensation inthe atmosphere.

Temperature has the opposite effect on thesecatalyzed reactions compared to other oxidationreactions: the formation of sulphate via thesemechanisms decreases as temperature decreases. Infact, the presence of the catalysts does not vary withtemperature, as is the case for the other oxidizingspecies found in the gaseous form, whilst the highactivation energies are affected by decreases intemperature.

Oxidation by Mn2� is inversely influenced bypH with respect to Fe3�: whereas in the case ofiron the reaction rate decreases as pH decreases inthe range 0-4, the opposite occurs for manganesein the range 0-3.

Fig. 6 shows the effect of pH on the various typesof oxidation. It can be observed that the dominantmechanism for the formation of sulphate at pH lowerthan 4-5 is that induced by hydrogen peroxide. Incontrast, when pH�5 oxidation by ozone is 10 timesfaster than that by H2O2. Oxidation catalyzed bymetals is important when the pH is high. Assuming awater content of 1 g m�3 within a cloud, the rate ofoxidation induced by H2O2 exceeds 100% an hour,whereas the rates of oxidation reactions catalyzed byiron and manganese are lower than 1% an hour whenpH �4.5.

The oxidizing power of nitrogen oxides withrespect to S(IV) differs depending on the speciesunder consideration. NO and HNO3 induce reactionswhich are too slow to contribute effectively tooxidation; HONO, on the other hand, has extremelylow atmospheric concentrations (1-8 ppb) and, despiteits high Henry’s constant (KH�49 mol l�1 bar�1), doesnot reach sufficient concentrations in solution torepresent an important oxidizing agent.

NO2, whose Henry’s coefficient is KH�1�10�2 moll�1 bar�1, by contrast, is a relatively insoluble gas;however, it has been shown (Schwartz, 1984) that thereaction rates of NO2 with HSO3

� and SO32� are

sufficiently high to make its oxidizing action important.Additionally, it has been shown that the reaction:

[105] 2NO2�HSO3�� 3H3O��2NO2

��SO42�

may be sufficiently fast at high pH levels (Lee andSchwartz, 1983; Lee, 1984).

Oxidation of nitrogenIn addition to oxidation from S(IV) to S(VI),

another possible oxidation process leading to theacidification of the water present in clouds or rain isthat induced by NO2 and NO and giving rise to nitricand nitrous acid:

[106] 2NO2(g)�H2O(l)�� 2H�(aq)�NO�

3(aq)�NO�2(aq)

[107] NO(g)�NO2(g)�H2O(l)�� 2H�(aq)�2NO�

2(aq)

However, these reactions do not contributesignificantly to acidification at the concentrations ofnitrogen oxides normally present in the atmosphere:the first of the two reactions (Schwartz, 1984) is veryslow, both due to the low solubility of NO2 and to thesecond order dependence of the reaction rate on theconcentration of NO2.

The formation mechanism for nitric acid in airremains primarily the same as the gas phase inducedby the free radical �OH:

[108] �OH �NO2 (�M)�� HONO2 (�M)

Analysing the composition of precipitation and thesurrounding air masses nevertheless shows thatoxidation in the aqueous phase also producessignificant quantities of nitric acid in the atmosphere(Lazrus et al., 1983; Misra et al., 1985).

A mixed gas-liquid mechanism for the productionof nitric acid in droplets has been proposed by Heikesand Thompson (1983):

[109] O3�NO2�� O2�NO3

[110] NO3�NO2 (�M)�� N2O5 (�M)

N2O5 generated in this way in the gas phase may beincorporated into droplets or onto the surface ofaqueous aerosol, and hydrolyze to HNO3:



rate

of

S(I

V)

oxid

atio

n (%

/h)

104

103

102

101

10�5

10�4

10�3

10�2

10�1

1

pH

O3, 50 ppb

H2O2, 1 ppb

HNO2, 1 ppb

NO2, 2 ppb

Fe, 3�10�7mol/l

Mn, 3�10�8mol/l

C, 10�2g/l

60 1 2 3 4 5

Fig. 6. Effect of pH on various sulphur(IV) oxidation reactions.

[111] N2O5�H2O�� 2H�(aq)�2NO�

3(aq)

However, a definitive evaluation of thecontribution of this mechanism to the concentration ofHNO3 in the liquid phase of the atmosphere is not yetpossible.

Special attention has recently been devoted toreactions in the aqueous phase involving speciesgenerated by photochemical mechanisms: inside cloudbodies there may be sufficient light for the productionof hydroxyl and hydroperoxide radicals.

Hydroxyl radicals react with nitrogen oxides toform nitrous and nitric acid:

[112] NO�2(aq)��OH(aq)�� NO2(aq)�OH�

(aq)

[113] NO(aq)��OH(aq)�� HNO2(aq)

[114] NO2(aq)��OH(aq)�� HNO3(aq)

Other potential reactions leading to the formationof nitric acid are those suggested by Heikes andThompson (1983) and Platt et al. (1984): the nitrateradical absorbed in solution is hydrolyzed to formnitric acid. This may occur starting from the radicalitself with the extraction of a hydrogen atom from anorganic compound:

[115] NO3��RH�� R��HNO3

The reaction with H2O2, however, does not appearto be important:

[116] HONO �H2O2�� NO3��H3O�

since it has been demonstrated (Lee, 1984) that thisreaction is too slow.

In the absence of clouds and in the presence ofsolar radiation, HNO3 is generated in the gas phase bythe reaction of NO2 with �OH radicals at a rate ofabout 20-30% per hour, whilst H2SO4 is generated bythis pathway far less effectively. By contrast, in theaqueous phase oxidation reactions involving H2O2, O3

and metal catalysts are able to produce H2SO4 insolution at rates of up to 100% an hour.

The chemico-physical properties of these two acidspecies are fundamentally different: nitric acid is farmore volatile and tends to remain in the atmosphere inthe gaseous form, whilst sulphuric acid has a relativelylow vapour pressure (�10�7 bar) and therefore tendsto be present inside particles. Both may react withbasic species. These reactions essentially occur withthe ammonia present to form ammonium nitrate andammonium sulphate:

[117] HNO3�NH3��

��NH4NO3

[118] H2SO4�2NH3��

��(NH4)2SO4

These two species are the main cause ofdiminishing visibility in the case of photochemical

smog. Ammonium nitrate is found in the solid form ifthe temperature is below that of deliquescence; ifrelative humidity is high, it is present in the aqueousphase. However, nitric acid may return to the gas phaseeven after forming ammonium salts.

Importance of aqueous phase equilibria in the formation of acid rain

Formed in the atmosphere by the mechanismsdescribed above in the aqueous and gas phases, theacids under examination may be deposited on theEarth’s surface by two main mechanisms: drydeposition and wet deposition. The distinction is madebased on the phase in which the pollutant is foundwhen it hits the surface: dry deposition involvesgaseous pollutants or small particles, whilst wetdeposition involves pollutants present in the dropletsinside fog, clouds and rain.

It should be specified that this classification onlyconsiders the transportation mechanism, and not thenature of the surface (whether or not it is wet or if itpresents a liquid film). Given the variable nature ofprecipitation it is difficult to make quantitativecalculations of the extent of wet deposition of apollutant species.

The deposition rate of a pollutant is givenapproximately by the product l�C, where C is theconcentration of the pollutant and l is the scavengingcoefficient proportional to the intensity of theprecipitation itself.

Dry deposition, in contrast, is characterized by thedeposition rate Vg, defined as:

[119] Vg�F�[S]

where F is the flux of species S towards the surfaceand [S] is the concentration of the same species at areference altitude h.

As far as wet deposition by precipitation isconcerned, the sulphate and nitrate deposited onexposed surfaces may be the result not only ofoxidation reactions in the aqueous phase, but also ofthe inclusion of aerosol particles containing thesespecies inside the cloud or of the scavenging ofaerosol particles beneath the cloud itself. If theseinclusion and scavenging mechanisms predominate,acidification will be determined by the parametersgoverning the formation of nitrates and sulphates inthe gas phase: solar radiation, the concentration ofNOx and hydrocarbons.

If formation inside the droplets by oxidation ofSO2 and NO2 prevails, however, the parametersdetermining acidity will be the concentrations ofoxidants such as H2O2 and O3.

Comparing the composition of interstitial aerosols (afraction of the atmospheric aerosol which, in the



presence of a cloud, remains the same and is notremoved by the water droplets present inside the cloud)and that of clouds provides information on the source ofacidity inside the clouds themselves. Daum andcolleagues (1983) have shown that for interstitialaerosols the ratio [H�]/[NH4

�] is less than 1, whereas theaerosol contained in clouds has a ratio greater than 1.

Similar conclusions have been reached by otherresearch groups (Lazrus et al., 1983; Harrison and Pio,1983): the higher acidity of the aerosols in cloud bodieswith respect to that of the surrounding atmosphereagrees well with the formation mechanisms of acidicspecies in the aqueous phase described above.

It can be concluded that the formation of sulphuricacid inside cloud aerosols provides the largestcontribution to the formation of sulphate in acid rain.

10.1.4 Depletion of stratosphericozone

The stratosphere is that region of the Earth’s atmospherewhich stretches from about 15 to 50 km above thetroposphere. The temperature is almost constant in thelowest layer, whilst it tends to increase gradually in theupper half. The temperature pattern in the stratosphereis regulated by variations in the concentration of ozoneinside this layer: by absorbing solar radiation in theultraviolet band ozone molecules convert solar energyinto kinetic energy, helping to heat the stratosphere.

The average global concentration of stratosphericozone varies with altitude; at an altitude of 15 km it isin the order of 0.5 parts per million (ppm) by volume,at around 35 km it increases up to 8 ppm and thendecreases to 3 ppm in the high stratosphere (45 km).The thickness of the ozone layer above a certaingeographic area or its average global value can beconventionally expressed in Dobson Units (DU), inother words with a height of 0.01 mm; the ozone wouldbe at this thickness if it were the only component of theatmosphere and if it were at a pressure of 1 bar and at atemperature of 0°C. On average the height of the ozonelayer may vary with latitude between 250 and 400 DU;its average global value is about 300 DU.

The concentration of ozone in the various regionsof the stratosphere is the result of a dynamic formationand destruction process. Ozone forms at an altitude ofabout 30 km, where ultraviolet solar radiation with awavelength shorter than 242 nm slowly dissociates theoxygen molecules into atomic oxygen:

[120] O2�hv�� O �O

The oxygen atoms rapidly combine with molecularoxygen to form ozone; in the presence of a third inertmolecule, M, the following reaction may take place:

[121] O �O2�M�� O3�M

In turn, the ozone molecules absorb the highenergy photons present in solar radiation, withwavelengths ranging from 240 to 320 nm. A directconsequence of this absorption is the dissociation ofozone into an oxygen molecule and an excited oxygenatom, O(1D), according to the reaction:

[122] O3�hv�� O2�O

This absorption process means that a significantportion of ultraviolet solar radiation does not reach theEarth’s surface: this has made the development of lifepossible on our planet. The photolysis of ozone is not agenuine destruction mechanism because virtually allthe oxygen atoms produced by this reaction combinerapidly with oxygen molecules to form ozone again.This mechanism entails the conversion of solar energyinto thermal energy, especially in the upper part of thestratosphere. The presence of ozone is thus the causeof the temperature inversion characterizing the upperbelt of the stratosphere.

The current distribution of stratospheric ozone ismostly determined by transportation processes. Ozoneis produced mainly at the tropics at an altitude ofbetween 25 and 35 km but, as a result of themovements of air masses, its highest concentrationsare found near the poles at an altitude of about 15 km.

Various destruction mechanisms contribute tobalancing out ozone formation processes (reactions[120] and [121]). An example is the reaction of ozonewith oxygen atoms to form molecular oxygen:

[123] O3�O�� O2�O2

The above reactions are known as Chapmanreactions and have formed the basis for the study ofstratospheric ozone. Chapman’s scheme onlyestimates the loss of ozone through naturaldestruction and does not consider the transportationof ozone onto the Earth’s surface which contributesa further 0.5%.

About 10% of destruction processes can beattributed to catalytic cycles involving species whichcontain hydrogen: free hydrogen atoms (H), hydroxyl(�OH) and hydroperoxide (HO2�) radicals give thesame results as reaction [123].

For example:

O �HO2�� OH �O2

HO��O3�� HO2��O211111111122

O3�O�� O2�O2

These species containing hydrogen are generatedby the reaction which normally occurs between watervapour and methane and the excited oxygen atoms O(1D) from the photolysis of ozone.



A significant contribution, about 70%, to theprocess of ozone destruction is supplied by a catalyticcycle involving NO and NO2, the main reactions ofwhich are:

O�NO2�� NO�O2

NO�O3�� NO2�O211111111122

O3�O�� O2�O2

The main source of NOx in the stratosphere is theoxidation of the nitrous oxide (N2O) produced bybacteria in the land and water. Although most nitrousoxide is converted into N2 and O by ultraviolet light,about 1% reacts with the excited oxygen atoms O (1D)generated by the action of ultraviolet radiation onozone to form nitrogen oxide and begin the cycle:

[124] O (1D) �N2O�� NO �NO

The NOx molecules, in turn, are removed by theformation of nitric acid (HNO3) in the reaction�OH-NO2 which occurs in the lowest layer of thestratosphere. Atmospheric currents transport the nitricacid towards the troposphere, where it is removed byprecipitation; in the same way, the nitrous oxide fromthe troposphere is transported upwards and destroyed.

The destruction of ozone may also be catalyzed bysubstances other than HOx e NOx, in particular bychlorine (Cl) and bromine (Br) atoms and theirrespective oxides (ClO and BrO).

The year 1930 saw the beginnings of the industrialproduction of chlorofluoromethanes andchlorofluoroethanes, known by the name freon, whichbecame widely used as liquid refrigerants, solventsand propellants for aerosol cans. In 1975, tworesearchers, F.S. Rowland and M.J. Molina, publishedan article stating that in the stratosphere freon couldinduce radical chain reactions, with negative effectsthe on the natural equilibrium of ozone. An example ofthe reactions taking place is as follows:• initial stage

[125] CF2Cl2�hv�� CF2Cl��Cl

• propagation stage

[126] O �ClO�� Cl��O2

[127] Cl��O3�� ClO��O2

In the initial stage, ultraviolet light causes thehomolytic cleavage of a C�Cl bond of the freon. Thechlorine radical may induce the propagation of a chainreaction leading to the destruction of ozone molecules.Again in 1976, a study by the National Academy ofSciences (NRC, 1976) confirmed Rowland andMolina’s predictions and in January 1978 the use offreon in aerosol cans was banned in the United States.

The main chlorine compounds present in thetroposphere (Table 6) are methyl chloride (CH3Cl), a

tiny proportion of which is of industrial origin, theman-made chlorofluoromethanes CFCl3 (CFC-11) andCF2Cl2 (CFC-12), and carbon tetrachloride (CCl4)generated by both natural and man-made sources. Lessimportant man-made sources include trichloroethylene(CCl2�CHCl) and the substances which havereplaced it, methyl chloroform (CH3CCl3) andtrichlorotrifluoroethane (C2F3Cl3), CFC-113. From the point of view of their effects on thestratosphere, the key species are thechlorofluorocarbons CFCl3 (CFC-11) and CF2Cl2(CFC-12) whose concentrations in the atmosphereincreased by 37% and 31% respectively between 1976and 1981. It is thought that the residence times ofCFC-11 and CFC-12 are about 60 and 110 yearsrespectively; measurements indicate a minimumlife-time for CFC-11 of 40 years.

A few years after their commercialization, thesechlorine compounds spread through the troposphere,and their concentration subsequently began to increaseslowly in the stratosphere as well.Chlorofluorocarbons (CFC-11, CFC-12, CFC-113) arehighly inert compounds in the troposphere and in thelower layers of the stratosphere, but when transportedto altitudes of 25-50 km they are decomposed byultraviolet radiation with wavelengths shorter thanabout 200 nm, with the resulting formation of chlorineatoms.

The chlorine atoms subsequently participate in thecatalytic cycle of ClOx. This cycle may be interruptedby the conversion of the highly reactive forms Cl andClO into less reactive forms which do not destroyozone. The chlorine atoms are deactivated by reactionwith methane to form HCl:

[128] Cl�CH4�� HCl�CH3



Table 6. Halogenated compounds presentin the troposphere

Compound NameChemicalformula

CFC-11 Trichlorofluoromethane CFCl3

CFC-12 Dichlorofluoromethane CF2Cl2

CFC-113 Trichlorotrifluoroethane C2F3Cl3

CFC-114 Dichlorotetrafluoroethane C2Cl2F4

CFC-115 Chloropentafluoroethane C2ClF5

Halon-1211 Bromochlorofluoromethane CF2ClBr

Halon-1301 Bromotrifluoromethane CF3Br

which acts as a temporary reservoir for the activechlorine species in the stratosphere. The chlorineatoms are regenerated by the reaction between HCland �OH radicals:

[129] �OH �HCl�� H2O �Cl

The destruction and regeneration processes of theactive chlorine species, ClOx, may occur several timesbefore the chlorine is completely removed from thestratosphere. The most important removal mechanismis the transportation of HCl from the stratosphere tothe upper troposphere, from which it is removed by theaction of rain, as occurs for the removal of nitric acidfrom the NOx cycle. The time scale from the point atwhich the chlorofluorocarbons are emitted at groundlevel to the point at which their chlorine atoms, in theform of HCl, are removed from the atmosphere by rainis in the order of several decades; variations in theemission fluxes of chlorofluorocarbons thus manifestthemselves in the stratosphere after many years.

Ozone destruction reactions may also be catalyzedby bromine atoms. It is thought that stratosphericbromine is 50 times more active than Cl in theprocesses leading to the destruction of ozone, and thatit is responsible for 20% of the ozone hole over theAntarctic and for a greater proportion of that over theArctic. The most important transportation vehicle forbromine is methyl bromide (CH3Br); the photolysis ofthis compound leads to the formation of the Br radical,responsible for the demolition of O3:

[130] CH3Br �hv�� CH3��Br

[131] Br �O3�� BrO��O2

[132] BrO��O�� Br �O2

Little is known about methyl bromide’s sources andcycle, unlike chlorofluorocarbons. This gas is used asa fumigant for the antiparassitic treatment of land andagricultural products and as a raw material forproducts of chemical synthesis. Together, natural sources(mainly marine biological activity; Lovelock, 1975) andman-made sources emit about 100 kt a year into theenvironment. Considering the emission fluxes and alife-time of about 2 years, the average global concentrationof methyl bromide is on the order of 10 ppt.

A particularly interesting aspect of the chemistryof stratospheric bromine is the possible synergicinteraction between its cycle and that of chlorine,through the reaction:

[133] BrO �ClO�� Br �Cl �O2

Yung and colleagues (1980) suggest that theinteraction between ClOx and BrOx may lead to anincrease in the ozone destruction process in the lowerlayers of the stratosphere.

In order to replace chlorofluorocarbons, othersubstances have been synthesized for use asrefrigerants and aerosol propellants:hydrochlorofluorocarbons (HCFCs) andhydrofluorocarbons (HFCs). The presence of at leastone hydrogen atom in each molecule means that theseare oxidized in the troposphere by reactions with �OHradicals:

[134] �OH �CHClxFy�� H2O �CClxFy�� products

The presence of hydrogen thus allows for a fasterdegradation of the substance in the atmosphere, andtherefore a lower environmental impact (HCFCs andHFCs do not reach the stratosphere).

The Ozone Depleting Potential (ODP) has beenintroduced to combat the harmful effects of thechlorofluorocarbons and the substances which havereplaced them. For any given halocarbon, the ODP isdefined as the ratio of stratospheric ozone destroyedby the emission of 1 kg of that compound to the ozonedestroyed by the emission of 1 kg of CFC-11.

A substance’s ozone depleting potential thusprovides a measure of the impact (in comparison tothat of CFC-11, considered the standard referencecompound with an ODP of 1.00) of the emission of 1kg of that substance in terms of the destruction ofstratospheric ozone. As such, the impact of a substanceon the ozone layer is given by its capacity to destroyozone and the quantity of total emissions. Table 7shows the ODP values and life-times in theatmosphere of the main compounds.



Table 7. ODP values and life-timesof the main halogenated compounds present in theatmosphere (WMO/United Nations Environment

Programme, 1994)

Compound ODPLife-time

in the atmosphere(years)

CFC-11 1.00 60

CFC-12 0.82 110

CFC-113 0.90 90

CFC-114 0.85 200

CFC-115 0.4 400

HCFC-22 0.04 13.3

HCFC-123 0.014 1.4

Halon-1211 5.1 20

Halon-1301 12 85

Other reactions may occur in the heterogeneousphase on the surface of the particles present in PolarStratospheric Clouds (PSCs), consisting of an icymixture of nitric acid and water, generated attemperatures below 195K (�78°C); these clouds,pearly white in colour, can frequently be seen over theAntarctic. HCl molecules aggregate onto the icecrystals, becoming part of the crystalline structure andleading to the formation of a solid phase catalyst. Thecatalyst reacts with chlorine nitrate to release chlorinein the gaseous state; this accumulates in thestratosphere and participates in the ozone destructionprocess:

[135] ClONO2�HCl�� Cl2�HNO3

The nitric acid remains bonded to the surface ofthe particles and therefore reacts with NO2, which isno longer able to neutralize ClO. In this way, the levelof active chlorine atoms is kept constant.

The low temperatures reached in polar Antarcticregions are associated with low pressure conditionsleading to the formation of a polar vortex which mixestroposphere and stratosphere whilst simultaneouslypreventing the entry of external air masses. The samelow temperatures encourage the formation of PSCs inthe Antarctic region, a phenomenon of minorimportance in the Arctic, where temperatures are about10-15°C higher.

It appears that aerosol particles containingsulphuric acid may present surface conditions thataccelerate the ozone destruction process. This wasshown clearly following the high emissions into thestratosphere of sulphur compounds in particulate andgaseous form by explosive volcanic eruptions (anexample is the eruption in 1991 of Mount Pinatubo inthe Philippines).

The release of large amounts of SO2 leads to theformation of aerosol through the reaction:

[136] �OH �SO2�� HSO3

which may considerably decrease theconcentrations of �OH radicals in the upper part of theatmosphere, with notable consequences. Thesubsequent reactions involving HSO3 are not fullyunderstood, but a potential further step might be:

[137] HSO3��OH�� H2SO4

Alternatively, two other reactions could besuggested:

[138] HSO3�O2�� HSO5

[139] HSO3�O2�� HO2��SO3

The oxidation of SO2 of volcanic origin in theatmosphere takes place very slowly. The sulphuric acidgenerated in this way may subsequently condense on

the water vapour present in small quantities in thestratosphere to form aerosol droplets.

Volcanoes are only one source of the gaseoussulphur present in the stratosphere; another, ofbiogenic origin, releases carbon disulphide. The levelsof carbon disulfide present in the stratosphere areextremely low, but it may be transferred there acrossthe tropopause, unlike other sulphur compounds suchas hydrogen sulphide and sulphur dioxide, apparentlytoo reactive to leave the troposphere. The highconcentrations of particulate matter containingsulphate (in the form of both SO4

2� and H2SO4) areshown in Fig. 7.

The diagram shows that carbonyl sulphide (OCS)plays an important role in the chemistry of the sulphurcompounds present in the atmosphere. It can beoxidized to sulphuric acid:

[140] OCS ��OH�� CO2�HS

[141] HS ��OH�� SO �H2

[142] SO ��OH�� SO2�H

It has been estimated that over 50% of the carbonylsulphide present in the atmosphere is man made and isemitted by combustion processes and the treatment offossil fuels.

Some models suggest that one of the mostsignificant consequences of ozone depletion will bethe cooling of the lower part of the stratosphere. It isthought that there is a positive feedback mechanism bywhich the loss of ozone cools the air, encouraging theformation of stratospheric polar clouds which in turn



H2SO4

SO42�

SO2

CS2

OCS

alti

tude

(km

)

40

0

10

20

30

concentration (molar fraction)10�12 10�910�1010�11

tropopause

similarlyH2S and(CH3)2S

Fig. 7. Concentration of the main sulphur compounds as a function of altitude.

contribute to the lowering of ozone levels. It is obviousthat the formation of PSCs is not limited to polarvortices, but may occur inside stratospheric jet streamsat temperate latitudes. All this suggests thatheterogeneous processes involving aerosols containingice particles, sulphuric acid and sulphur compoundsare important components of the chemistry of thestratosphere.

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Ivo AllegriniConsiglio Nazionale delle Ricerche

Istituto sull’Inquinamento AtmosfericoMonterotondo, Roma, Italy



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10 ENVIRONMENTAL TECHNOLOGIES - treccani.it · The term nitrogen oxides is generally used to describe a mixture of nitrogen monoxide (95% of the total), nitrogen dioxide and traces