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10
ENVIRONMENTAL
TECHNOLOGIES
10.1
Atmospheric chemistry
10.1.1 Introduction
In recent years, growing concern about environmental
issues has focused attention on the quality of
atmospheric air, both in the troposphere, in other
words from the ground up to an altitude of about 10
km, and in the stratosphere, due to the well-known and
worrying phenomenon of the decrease in the
concentration of ozone and the consequent thinning of
the layer protecting the Earth from ultraviolet
radiation. This interest has resulted in the creation of
specific environmental regulations, which made their
first appearance in the middle of the Twentieth
century.
In Europe, this legislation sets out limiting values
and objectives to be reached by a certain date in the
form of Directives (among the Directives currently in
force it is worth recalling 96/62/EC, 99/30/EC,
2000/69/EC and 2002/3/EC). Their adoption by each
member state means that these Directives become state
law. Additionally, on an international level,
commissions have been established to deal with
specific problems and to draft protocols (the most
important include the Geneva Protocol of 1979, the
Montreal Protocol of 1987 and the Kyoto Protocol of
1997, with their subsequent amendments) subject to
ratification by individual countries. An attempt has
thus been made to take account of the impossibility of
restricting to a single state or continent problems
which inevitably affect neighbouring states (so-called
transboundary pollution) or which, in the case of
extremely complex phenomena such as climate change
or the depletion of stratospheric ozone, affect the
entire planet on a global scale. The latter two
phenomena have been and remain (especially the first)
an object of study by the scientific community and of
interest to the authorities charged with managing the
environment. The depletion of stratospheric ozone will
VOLUME III / NEW DEVELOPMENTS: ENERGY, TRANSPORT, SUSTAINABILITY
be dealt with below, as part of a brief overview of the
complex interactions between the atmospheric ‘vessel’
and the compounds of natural and man-made origin
emitted into it, thus creating a complex system. This
system can only be understood by taking into account
both the chemistry of the aforementioned compounds
and the meteo-climatic factors which can alter their
concentrations, for example by preventing their mixing
and consequently their dispersion (episodes of thermal
inversion are in fact closely correlated with cases of
extreme pollution).
There are four basic processes underlying pollution
phenomena: emission, transformation, diffusion and
transportation, and deposition.
The first two processes essentially concern those
pollutants described as ‘primary’, in other words those
emitted directly as such, which may then undergo
transformations and give rise to ‘secondary’ pollutants
(such as nitrogen dioxide and ozone). These two
classes of pollutants share their subsequent ‘itinerary’,
which depends on the dynamic behaviour of the lower
layers of the atmosphere, as far as transportation and
diffusion mechanisms are concerned, and is heavily
conditioned by this and by other meteo-climatic
parameters (such as relative humidity), as far as sink
processes through dry and wet deposition are
concerned.
This chapter focuses on the chemistry of the
compounds of greatest environmental interest
(nitrogen oxides, sulphur dioxide, non-methane
organic compounds, ozone). Specifically, the reactions
characterizing these in both the gas phase
(homogeneous reactions) and in the aqueous phase
(heterogeneous reactions) will be examined: the
presence of water in the atmosphere, in different states
of aggregation, is important for general considerations
on the behaviour of these species after their release
into the atmosphere.
915
ENVIRONMENTAL TECHNOLOGIES
An in-depth understanding of all the processes
governing pollution phenomena is essential both for
the ability to predict how these will evolve in the short
and long term, and to identify and implement
appropriate control strategies.
(or nitrous oxide), nitrogen trioxide, tetroxide and
pentoxide. Nitrogen monoxide is a colourless and
odourless primary pollutant which forms from
combustion at about 1,200°C, whilst nitrogen dioxide
is a reddish pollutant with a strong and pungent smell.
Nitrogen dioxide is a secondary pollutant since it does
not form directly from combustion (except in a
percentage of 4-5%).
The sequence of reactions leading to the formation
of nitrogen oxides is described below.
Combustion produces NO (primary pollutant)
through the reaction:
10.1.2 Atmospheric chemistry
in the gas phase
The apparent stability of the atmosphere derives from
the fact that it is in a stationary state; this situation is
due to the relative constancy of inputs and outputs
which on average balance each other out. Since most
sources and sinks of the gases in the atmosphere are
the result of chemical reactions, an understanding of
reaction rates in the gas phase is extremely important.
However, the study of chemical kinetics concentrates
not only on reaction rates but also on determining the
precise reaction mechanism.
This paragraph provides a brief survey of the
reactions of the main pollutants in the gas phase.
Specifically, Fig. 1 shows the reactions induced by the
chemical species most heavily involved in the
formation of acidic compounds in the atmosphere.
[1]
䉴
which in turn can produce NO2 (secondary pollutant)
via the (non-photochemical) thermal oxidation
reaction:
[2]
2NO ⫹O2⫺ 2NO2
䉴
However, reaction [2] is of minor importance since
it is too slow at the concentrations of NO normally
present in the atmosphere. The reaction kinetics is of
the second order with respect to NO and of the first
order with respect to O2; therefore, the kinetic
equation is R⫽k [O2] [NO]2, where R is the reaction
rate 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 ultraviolet
radiation and tropospheric ozone.
Nitrogen oxides
The term nitrogen oxides is generally used to
describe a mixture of nitrogen monoxide (95% of the
total), nitrogen dioxide and traces of dinitrogen oxide
Fig. 1. Some reactions
which occur
in the atmosphere induced
by NOx, SO2 and NH3
(the species most involved
in the formation
of acid compounds
in the atmosphere
and in their subsequent
salification).
N2⫹O2⫺ 2NO
O3
O3
NO
aerosol
OH
HNO3
NO2
hn
nitrate aerosol
NH3
emission
dry deposition
OH
wet deposition
NH3
SO2
H2SO4
sulphate aerosol
H2O2
emission
dry deposition
wet deposition
NHO3
ammonium aerosol
NH3
H2SO4
emission
916
dry deposition
wet deposition
ENCYCLOPAEDIA OF HYDROCARBONS
ATMOSPHERIC CHEMISTRY
The photolysis of NO2 produces ozone through the
reactions:
[3]
[4]
NO2 ⫹hv (l ⭐430 nm)⫺ NO ⫹O(3P)
O(3P) ⫹O2⫺ O3
(k ⫽6.0 ⭈10⫺34 (T/300) ⫺2,3 cm6 molecules⫺1s⫺1)
䉴
䉴
where hv is the light radiation and l is the wave length.
The reaction between ozone and nitrogen
monoxide then produces nitrogen dioxide:
O3 ⫹NO⫺ NO2 ⫹O2
(k ⫽1.8 ⭈10⫺14 cm3 molecules⫺1s⫺1)
There is not therefore a net loss of ozone, since the
titration reaction causes it to be destroyed whilst the
NO2 photolysis reaction simultaneously causes
reemission; it can thus be deduced that the three
species are in fact involved in a photostationary
equilibrium. The rate constant in decomposition
processes by photolysis of a trace species present in
the atmosphere, whose dimensions are the inverse of
time (s⫺1), can be generically defined as:
[5]
[6]
䉴
冮
J ⫽ s(l)f(l)j(l)dl
where s(l) is the absorption cross-section expressed in
cm2 molecules⫺1, which is characteristic of each
chemical species, f(l) is the quantum yield of the
photolysis reaction, and finally j(l) is the actinic flux
(number of photons⭈cm⫺2⭈s⫺1, generally used instead
of I⬘i ⫽I⬘r⫹I⬘s⫹I⬘d); in other words it is the product of the
components of the solar radiation incident (I⬘i ) directly
on the air masses, consisting of reflected radiation (I⬘r ),
scattered radiation (I⬘s ), and finally that which reaches
the Earth’s surface directly (I⬘d ).
Assuming the absence of organic compounds in
the atmosphere, the relationship between the
concentrations of O3, NO and NO2 at a generic time t
in a mass of air is a constant given by the ratio of JNO2
(the rate constant of NO2 photolysis) to kNO (the rate
constant of reaction [5] between NO and ozone):
[7]
JNO2[NO2]
[O3]⫽ 11112
kNO[NO]
The constant JNO2 varies depending on the angle of
the luminous solar radiation, and so the relationship
between the concentrations changes over the course of
a day. The average value of JNO2 is 0.75⭈10⫺2 s⫺1 or
0.533 min⫺1.
The presence of RO⭈2 and HO⭈2 radicals, which
form mainly during the hot season, promotes reactions
with nitrogen monoxide, in other words:
[8]
RO⭈⫹NO⫺
NO2 ⫹RO
2
(if R ⫽CH3CH2CH2,
k ⫽7.6 ⭈10⫺12 cm3 molecules⫺1s⫺1)
䉴
VOLUME III / NEW DEVELOPMENTS: ENERGY, TRANSPORT, SUSTAINABILITY
[9]
HO2⭈⫹NO⫺ HO ⫹NO2
(k ⫽8.3 ⭈10⫺12 cm3 molecules⫺1s⫺1)
䉴
As a result, under these conditions, reaction [5] is
hindered and O3, being unable to react with NO,
accumulates in the lower levels of the atmosphere.
In the presence of NO2, high concentrations of RO⭈2
and HO⭈2 free radicals give rise to alkyl nitrate
hydroperoxides, PeroxyAcetyl Nitrates (PAN) and
PeroxyBenzoyl Nitrates (PBN).
PAN is a gas which tends to accumulate
persistently in the coldest parts of the troposphere; it
then spreads to warmer areas, leading to the formation
of free radicals and NO2 since the dissociation
constant of PAN depends strongly on temperature.
World Health Organization (WHO) guidelines state
that the annual averages of NOx in European cities are
around 40 mg/m3 and range from 20 to 90 mg/m3 in
industrialized countries; the base level is between 1
and 9 mg/m3.
Nitrogen dioxide is more toxic than nitrogen
monoxide. For this reason, NO2 rather than NO is
monitored by law; values of 13 ppm cause irritation to
the mucous membranes of the nose and eyes; exposure
to concentrations of 560 mg/m3 for 30 minutes causes
pulmonary problems. The WHO recommends that
hourly mean concentrations of 200 mg/m3 and an
average annual limit value of 40 mg/m3 should not be
exceeded.
As already seen, NO2 is a pollutant which is
mainly generated indirectly from the NO emitted by
the combustion of fuels used in road transportation
and is thus described as a ‘mobile source pollutant’. It
is found mainly in urban areas with a high traffic
density, coinciding with the opening and closing times
of workplaces and schools, and is also present in large
car parks. Another source of this substance is
combustion processes in civil and industrial plants.
As well as being harmful to human health, nitrogen
dioxide acidifies rain, degrades man-made objects,
corrodes metals and damages vegetation.
During the day, NO2 is oxidized to nitric acid by
reaction with the ⭈OH radical:
[10]
NO2 ⫹⭈OH ⫹M⫺ HNO3 ⫹M
(k ⫽1.1⭈10⫺11 cm3 molecules⫺1s⫺1)
䉴
where M is a third ‘body’ or molecule (typically N2 or
O2) that has the role of removing the excess energy
through collision and so influences the kinetics of the
reaction itself.
This reaction is slow compared to the NO-NO2
exchange: the average life-time of NOx is typically 1-2
days in the low troposphere at middle latitudes. The
nitric acid is then deposited on the Earth’s surface
through various mechanisms.
917
ENVIRONMENTAL TECHNOLOGIES
The NO3 radical is particularly important in the
chemistry of organic compounds during the night,
when its average life-time increases since it does not
undergo photodissociation (as occurs during the day)
and is thus able to react with Volatile Organic
Compounds (VOCs), oxidizing them rapidly. Its action
is similar to that of ⭈OH during the daytime.
The reactions which take place can be grouped as
follows:
• formation reactions:
[11]
[12]
[13]
•
䉴
䉴
NO2 ⫹NO3 ⫹M⫺
⫺N2O5 ⫹M
䉳
䉴
For this reaction, the equilibrium constant reported in
the literature varies by a factor of 1.9 at ambient
temperature (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:
䉴
NO3 ⫹hv (l ⭐640 nm)⫺ NO2 ⫹O(3P)
NO3 ⫹hv (585 ⭐ l ⭐640 nm)⫺ NO ⫹O2
䉴
䉴
reactions with organic compounds
[16]
NO3 ⫹RH⫺ HNO3 ⫹R
[17]
NO3 ⫹RCHO⫺ HNO3 ⫹RCO
[18]
NO3 ⫹C⫽C⫺ ⫺C(ONO2)⫺C⫺
䉴
䉴
䉴
Addition reactions to alkenes (especially biogenic
isoprene and monoterpenes) are much faster than
those which occur with alkanes. The reaction of the
nitrate with olefins also forms peroxide radicals (HO⭈2
⫹RO⭈)
2 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 with
hydrocarbons accounts for about 15% of the nitric acid
in the atmosphere. Considering the average
concentrations of these organic compounds in the
atmosphere (in typically polluted urban atmospheres
alkanes are present in concentrations of about 100 ppb
whilst formaldehyde and acetaldehyde are present
respectively in concentrations of 20 ppb and 10 ppb)
and the rate constants, and assuming that the
concentration of NO3 is 100 ppt, the net overall rate of
formation is about 0.3 ppb h⫺1. Making similar
assumptions for an average concentration of NO2 of
50 ppb and ⭈OH concentrations of 1⭈106 cm⫺3 (typical
of a moderately polluted atmosphere) the rate of
formation is 2 ppb h⫺1.
The nitrate is thus a sink for NOx during the night,
and may in turn be removed by various mechanisms
involving other sinks such as reactions with organic
compounds or deposition on aerosols or on the ground
(direct sink). It has in fact been shown that the average
life-time of NO3 depends on the relative humidity
decreasing rapidly to less than 10 minutes when
918
[19]
[20]
photodissociation reactions
[14]
[15]
•
NO ⫹O3⫺ NO2 ⫹O2
(k ⫽1.8 ⭈10⫺14 cm3 molecules⫺1s⫺1)
NO2 ⫹O3⫺ NO3 ⫹O2
(k ⫽3.2 ⭈10⫺17 cm3 molecules⫺1s⫺1)
NO2 ⫹⭈OH⫺ HNO3
(k ⫽1.1⭈10⫺11 cm3 molecules⫺1s⫺1)
relative humidity is 50% (Platt et al., 1984). This is
probably due to interaction with the water found on the
surfaces of the particles present in the environment.
In part, this may also be explained by the reaction
of nitrogen pentoxide (N2O5) which represents an
indirect sink for NO3 and which may react in the
presence of water, shifting the equilibrium:
N2O5 ⫹H2O⫺ 2 HNO3
(k ⭐1.3 ⭈10⫺21 cm3 molecules⫺1s⫺1)
䉴
Homogeneous hydrolysis thus represents a sink for
NO3 and at the same time a source of nitric acid:
although the rate constant seems fairly low, this reaction
contributes significantly to the formation of nitric acid
in the homogeneous phase (0.3 ppb h⫺1 at 50% of
relative humidity) as well as in the heterogeneous
phase, and in this case is significantly faster.
Essentially, the main formation paths for nitric acid
are the reaction between NO2 and ⭈OH, the reaction
between nitrate and organic compounds and the
hydrolysis of N2O5. Nitric acid has a fairly long
life-time, and may thus be the terminal of various
chain reactions in the troposphere. It absorbs weakly
in the actinic region and thus does not undergo
photolysis, however it may undergo deposition (dry
and wet) and react with the ⭈OH radical and (albeit
slowly) with ammonia:
[21]
[22]
HNO3 ⫹⭈OH⫺ H2O ⫹NO3
HNO3 ⫹NH3⫺ NH4NO3
䉴
䉴
Ammonium nitrate is thus in equilibrium with the
two species in the gas phase, and this equilibrium
exists for both ammonium in the solid phase
(relative humidity⬍62%) and in solution (relative
humidity⬎62%).
Sulphur dioxide (SO2)
Sulphur dioxide (SO2) is a colourless gas with an
acrid, pungent smell. Emissions of this gas are mainly
due to the use of solid and liquid fuels and are directly
correlated with their sulphur content:
[23]
S ⫹O2⫺ SO2
䉴
This is, therefore, a typical pollutant of industrial
and urban areas, in the latter particularly during the
winter (due to domestic heating). Natural
concentrations of SO2 are less than 5 mg/m3; annual
ENCYCLOPAEDIA OF HYDROCARBONS
ATMOSPHERIC CHEMISTRY
averages are lower than 50 mg/m3; and daily averages
rarely exceed 125 mg/m3.
Sulphur dioxide is used as a ‘tracer’, in other
words a global indicator of atmospheric pollution, due
to its chemical stability in the atmosphere.
Most SO2 undergoes chemical transformations
before it reaches the ground; it is oxidized to SO3,
followed by hydrolysis to H2SO4, particles of which
absorb further SO2, NH3 and traces of metals to form a
particulate aerosol which, depending on weather
conditions, may be transported for hundreds of
kilometres and reach the ground in the form of acid rain.
The oxidation reaction of SO2:
[24]
2SO2 ⫹O2⫺ 2SO3
䉴
has such a low rate in the absence of catalysts that it
can be completely disregarded as a source of SO3; the
same can be said for photooxidation as a reaction
mechanism, since if every SO*2 molecule in an excited
state was oxidized by reaction with O2 or other species,
the average life-time of SO2 in the low troposphere
would be 52 minutes, which is, in fact, not the case.
In any case, the only rapid gas phase process which
is efficient enough to account for most of the sulphuric
acid present in the aerosols formed by gas phase
processes is the reaction of SO2 with the ⭈OH radical:
[25]
SO2 ⫹⭈OH ⫹M⫺ HOSO⭈⫹M
2
(k ⫽1.1⭈10⫺12 cm3 molecules⫺1s⫺1)
䉴
It is known that a significant fraction of the
HOSO⭈2 radical is eventually transformed into
sulphuric acid, but the rate of the reaction and its
intermediate products are not well known; as such, the
reaction is generically described as:
[26]
HOSO⭈⫺
⫺ ⫺ H2SO4
2
䉴
䉴
䉴
In any case, a mechanism by which ⭈OH is
regenerated has been suggested:
[27]
[28]
HOSO2 ⫹O2⫺ HO2⭈⫹SO3
HO⭈⫹NO⫺
NO2 ⫹OH
2
䉴
䉴
This mechanism was proposed by Calvert and
Stockwell (1984), on the basis of the experimental
evidence that in a photooxidant mixture of HNO2, NO,
NO2 and CO, even the addition of substantial amounts
of SO2 does not influence the concentration of ⭈OH.
The oxidation of SO2 by this mechanism averaged
out over 24 hours is 16.4%; during the winter the rate
is lower due to the lower concentration of ⭈OH.
Other oxidation reactions induced by other oxidizing
species, such as O (3P), HO⭈2 and CH3O⭈,
2 are
characterized by lower rate constants than the reaction
with ⭈OH, which remains the main oxidation reaction
induced by SO2 (the rate constants are k⫽5.7⭈10⫺14,
k⬍1⭈10⫺18, and k⬍1⭈10⫺18 respectively).
VOLUME III / NEW DEVELOPMENTS: ENERGY, TRANSPORT, SUSTAINABILITY
Levels of SO2 are generally far higher than those of
SO3 since the latter, in contact with water vapour,
leads to the following reaction, also encouraged by the
presence of particulate matter and solar radiation:
[29]
SO3⫹H2O⫺ H2SO4
(k ⫽9.1⭈10⫺13 cm3 molecules⫺1s⫺1)
䉴
Naturally, other fundamentally important oxidation
reactions of SO2 to sulphuric acid are those which take
place in the aqueous phase inside the water droplets
present in the atmosphere; these will be dealt with in
the 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 with
respect to attack by ⭈OH. This means that it can be
transported far from its source before it is destroyed.
By marked contrast, terpenes are extremely reactive
and as a result have short life-times. A broad variety of
hydrocarbons, the so-called Biogenic Non-Methane
Hydrocarbons (BNHC), are emitted by natural
sources: these are unsaturated organic compounds
emitted mainly by plants, such as isoprene
and some monoterpenes like a-pinene, b-pinene,
d-limonene, etc.
In towns the air may contain variable
concentrations (even above 1.0-2.0 mg/m3) of
hydrocarbons other than methane, whilst methane
itself may exceed 1.5 mg/m3.
The most important reactions involving these
species (RH) are those which lead to the formation of
ozone:
[30]
RH⫹⭈OH ⫹O2⫺ RO⭈(⫹H
2O)
2
[31]
RO ⫹NO2
RO⭈⫹NO⫺
2
[32]
NO2 ⫹O2 ⫹hv⫺ O3 ⫹NO
䉴
䉴
䉴
As such, the amount of ozone produced per mole
of hydrocarbon depends on its concentration and its
reactivity with ⭈OH.
The relative reactivity of species j in a mixture of i
hydrocarbons (HC) can thus be expressed as:
kj [HC]j
[33] Rj ⫽ 111133
冱ki[HC]i
i
Taking account of the typical abundances of these
compounds and their reactivity, it is possible
to calculate the rates of degradation. For example,
Table 1 shows the removal rates of some of the most
abundant hydrocarbons in the atmosphere.
In terms of the rate of degradation, alkenes have
the highest potential for forming ozone, followed by
aromatic compounds, whilst alkanes have the lowest
formation potential.
919
ENVIRONMENTAL TECHNOLOGIES
Table 1. Kinetic data on the removal of hydrocarbons present in the atmosphere by ⭈OH radicals
k 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
COMPOUND
The oxidation of hydrocarbons is strictly correlated
with the photostationary state NOx⫺O3 described by
the Leighton relationship: the oxidation of VOCs
(Fig. 2) by ⭈OH produces the radical RO⭈2 which
converts NO to NO2 through the reaction already
described for NOx without involving the simultaneous
removal of O3. As a result, the concentration of ozone
increases, forming other ⭈OH radicals, thus increasing
the rate at which VOCs are oxidized.
The maximum production of O3 requires an
appropriate ratio of NOx to VOC concentrations: a lack
of NOx results in an insufficient production of ⭈OH to
induce the oxidation of VOCs, whilst a lack of VOCs
makes it impossible to reach the concentrations of
RO2⭈ needed to alter the photostationary state
significantly.
The non-linearity of ozone formation has been
demonstrated experimentally by Kelly and Gunst (1990).
Surveying the main reactions of the various classes
of organic compounds, it can be seen that the main
reactions are those with ⭈OH and NO3, whose
constants for the various hydrocarbons are reported in
Table 2.
The chemistry of the higher term alkanes follows
the same mechanism as methane: the attack by ⭈OH
takes place preferentially to form the most stable alkyl
radical, therefore tertiary and secondary hydrogens are
those which react most easily with ⭈OH. The chemistry
of anthropogenic alkenes follows a similar mechanism
to that of biogenic isoprene: the initial reaction is the
addition of ⭈OH followed by the addition of O2 to a
hydroxy-substituted alkylperoxy radical, which in turn
reacts with either NO or HO⭈2 depending on whether
the concentrations of NO are high or low.
The photochemical oxidation of carbonyl
compounds leads to the production of peroxyacetyl
radicals and the following reactions give rise to the
formation 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)
䉴
䉴
VOC ⫽RCH3
ⴢOH, O2, M
RO2ⴢ ⫽RCH2O2
NO
via aldehydes
NO2
⫹
ROⴢ ⫽RCH2O
hn
NO
O
O2
O3
Fig. 2. Schematic of the oxidation of organic
compounds in the troposphere.
920
O2
O2
Table 2. Rate constants
for VOC oxidation reactions
R⫹CHO
CONCENTRATION kOH(1012cm3 kNO3(1016cm3
molec.⫺1s⫺1) molec.⫺1s⫺1)
(ppb carbon)
hn
COMPOUND
RCHO
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
ENCYCLOPAEDIA OF HYDROCARBONS
ATMOSPHERIC CHEMISTRY
[36]
CH3C(O)O2NO2 ⫹M⫺ CH3C(O)O2 ⫹NO2⫹M
(k ⫽1.8 ⭈10⫺4 cm3 molecules⫺1s⫺1)
䉴
The formation of PAN is a process which ends the
propagation of the reaction chain, and which competes
with 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)
䉴
䉴
䉴
䉴
The formation of PAN is encouraged by low
temperatures and pressures. Thermal decomposition is
the most important destruction path for PAN near the
Earth’s surface, whilst at altitudes above 7 km it reacts
with ⭈OH:
[41]
CH3C(O)O2NO2⫹⭈OH⫺ products
(k ⫽1.1⭈10⫺13cm3 molecules⫺1s⫺1)
䉴
Therefore, if the PAN formed rises rapidly to
height atmosphere, its life-time increases and it may
represent a source of NOx through long-distance
transportation mechanisms.
Ozone and photochemical smog
For a photochemical smog process to be triggered,
sunlight, nitrogen oxides and volatile organic
compounds must be present; additionally, the process
is favoured by a high atmospheric temperature. Since
nitrogen oxides and volatile organic compounds are
among the main components of emissions in urban
areas, towns located in geographical areas
characterized by intense solar radiation and high
temperatures (such as those in the Mediterranean) are
ideal candidates for episodes of intense photochemical
pollution. The knowledge necessary for understanding
secondary pollution events thus concerns the chemical
and chemico-physical transformation processes
undergone by pollutants, dynamic processes in the
lower atmosphere (atmospheric stability, direction and
intensity of the wind) and the intensity of solar radiation.
Ozone is a photochemical oxidant similar to PAN,
nitrogen dioxide and hydrogen peroxide which are
essentially secondary pollutants formed in the
troposphere by chemical reactions starting from
primary pollutants (basically VOCs and nitrogen
oxides which are therefore also described as
precursors) in the presence of solar radiation. It is
naturally present in the troposphere at concentrations
ranging from 20 to 80 ppb.
VOLUME III / NEW DEVELOPMENTS: ENERGY, TRANSPORT, SUSTAINABILITY
Ozone has a characteristic smell and may cause
severe irritation to the respiratory system and the eyes
at concentrations exceeding 200 ppb. It is also a cause
of the oxidative degradation of some non-biological
materials, especially elastomers, textile fibres and
dyes.
In fact, the presence of only six electrons in the
valence shell of oxygen gives it electrophilic
properties and therefore the tendency to remove
electrons from other species or to share them. It is
characterized by a redox potential of 2.07 V in an
aqueous system.
The formation reactions of ozone by nitrogen
oxides and hydrocarbons have already been reported
extensively in the chapter devoted to those pollutants.
In the lower atmosphere, ozone forms from the
reaction of atmospheric oxygen with the atomic
oxygen produced by the photolysis of nitrogen
dioxide; the ozone formed is in turn removed by
nitrogen monoxide, with the new formation of NO2.
In unpolluted atmospheres, where other chemical
species are not present in appreciable quantities, this
series of reactions forms a cycle (the photostationary
ozone cycle) and there is no possibility for
photochemical pollution. The fundamental step
towards the atmospheric enrichment of ozone and
other photooxidant species (in other words oxidizing
chemical species formed by chemical reactions that
occur only in the presence of light) is the formation of
NO2 by alternative pathways which do not entail the
removal of ozone. Identifying the formation pathways
of NO2 thus represents the key to understanding
photochemical oxidation processes.
The main alternative pathway for the formation of
NO2 is the oxidation of NO by peroxide radicals
(RO⭈).
2 These free radicals originate from the
degradation of volatile hydrocarbon molecules (RH)
and their subsequent reaction with atmospheric
oxygen. The attack on volatile hydrocarbons is due to
the presence in the atmosphere of other free radicals,
⭈OH hydroxyl radicals: the processes which generate
hydroxyl radicals are thus fundamental in triggering
photochemical pollution processes. The production of
⭈OH radicals is essentially a photochemical process
and the main precursors are nitrous acid, formaldehyde
and ozone itself. Ozone is therefore not only the most
quantitatively important product of photochemical
pollution processes, but also part of the ‘fuel’ which
activates the process.
In fact, ozone, like nitrogen dioxide, undergoes
photolysis and given the energy of the O2⫺O bond
which is only 101 kJ mol⫺1 (in other words in the
order of 1 eV per molecule) and since the energy E of
a photon is linked to frequency n by the relationship
E⫽hn, fairly low frequency radiation is needed (but in
921
ENVIRONMENTAL TECHNOLOGIES
any case over 2.53⭈1014 Hz, in other words with a
wavelength shorter than 1.185 nm) to cleave the bond;
most of the spectrum of solar radiation thus has
sufficient energy to cleave the O2⫺O bond:
[42]
O3 (1A1) ⫹hv (l⬍1,180 nm)⫺ O(3P) ⫹O2(3S g⫺)
(J ⫽4.2 ⭈10⫺4 s⫺1)
䉴
At shorter wavelengths, the photon’s excess energy
may be converted into the electronic excitation of the
products:
[43]
O3 (1A1)⫹hv (l⬍310 nm)⫺ O(1D) ⫹O2(1Dg)
(J ⫽2.9 ⭈10⫺5 s⫺1)
䉴
Electronically excited molecular oxygen (1Dg) may
be another potential oxidizing agent in the
troposphere, especially for unsaturated hydrocarbons:
the rate constants for reactions with these compounds
are in the order of 10⫺18 cm3 molecules⫺1s⫺1.
Comparing this with the rate constants of the reactions
of alkenes with ⭈OH and O3 (on the order of 10⫺12, as
reported in Table 1, and 10⫺18 cm3 molecules⫺1s⫺1
respectively) leads to the conclusion that, to have a
significant effect on tropospheric chemistry, the
concentrations of O2 (1Dg) would need to be of the
same order of magnitude as ozone or greater. In fact,
concentrations of this species are around 10 ppt.
Since the transition:
[44]
O
O(1D) ⫹H2O⫺ 2⭈OH
(k ⫽2.2 ⭈10⫺10 cm3 molecules⫺1s⫺1)
䉴
This reaction is extremely fast and occurs in
competition 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)
922
R(⭈OH, O3) ⫽4.65 ⭈105 molecules cm⫺3 s⫺1
HONO ⫹hv(l ⬍320 nm)⫺ ⭈OH ⫹NO
(J ⫽1.8 ⭈10⫺3 s⫺1)
䉴
H2O2 ⫹hv(l ⬍360 nm)⫺ 2⭈OH
(J ⫽6.9 ⭈10⫺6 s⫺1)
In the case of formaldehyde, the predominant
mechanism when 300⭐l ⭐320 nm is:
[49]
[50]
䉴
HCHO ⫹hv ⫺ H ⫹CHO (J ⫽1.7 ⭈10⫺5s⫺1)
䉴
Subsequent oxidation leads to the formation of the
⭈OH radical:
[51]
H ⫹O2 ⫹M⫺ HO⭈⫹M
2
(k ⫽7.0 ⭈10⫺13 cm3 molecules⫺1s⫺1)
[52]
CHO ⫹O2⫺ HO⭈⫹CO
2
(k ⫽6.0 ⭈10⫺12 cm3 molecules⫺1s⫺1)
[53]
HO⭈⫹NO⫺
⭈OH ⫹NO2
2
(k ⫽8.0 ⭈10⫺12 cm3 molecules⫺1s⫺1)
䉴
䉴
䉴
At wavelengths above 340 nm there is a
preferential dissociation into relatively stable products:
[54]
HCHO ⫹hv⫺ H2 ⫹CO (J ⫽4.3 ⭈10⫺5s⫺1)
䉴
Crutzen (1988) has estimated that on average
50-60% of the formaldehyde follows the mechanism
above, 20-25% follows the mechanism leading to the
formation of CHO and H and subsequent oxidation
reactions, and 20-30% reacts directly with ⭈OH
according to the reaction:
[55]
䉴
In an atmosphere characterized by a relative
humidity of 50% and a temperature of 298K, about
10% of the O (1D) produced reacts with water to form
hydroxyl radicals.
At middle latitudes, the main source of ⭈OH
radicals is ozone, given the high concentrations of this
species which is about 40 ppb or 9.84⭈1011 molecules
cm⫺3.
In fact, the rate of production of ⭈OH radicals from
the photolysis of ozone at middle latitudes is:
[47]
[48]
(1D)⫺䉴 O (3P)
is impossible, as this is a spin-forbidden transition, O
(1D) may undergo thermal degradation through
collisional energy transfer or react with other species,
like CH4 or H2O, extracting a proton from them and
thus giving rise to ⭈OH radicals:
[45]
Other sources of ⭈OH radicals are the photolysis of
nitrous acid, hydrogen peroxide and formaldehyde.
Nitrous acid and formaldehyde are precursors of
⭈OH radicals, but in their turn have formation
pathways which are essentially secondary, starting
from species involved in photochemical processes
(nitrogen dioxide for nitrous acid and hydrocarbons
and radicals or ozone for formaldehyde).
The formation reactions for ⭈OH radicals starting
from nitrous acid and hydrogen peroxide are as follows:
HCHO ⫹⭈OH⫺ HCO ⫹H2O
(k ⫽1.0 ⭈10⫺11 cm3 molecules⫺1s⫺1)
䉴
The ⭈OH radical may also give rise to the oxidation
of CO:
[56]
CO ⫹⭈OH⫺ CO2 ⫹H
(k ⫽2.0 ⭈10⫺13 cm3 molecules⫺1s⫺1)
[57]
H ⫹O2 ⫹M⫺ HO⭈⫹M
2
(k ⫽7.0 ⭈10⫺13 cm3 molecules⫺1s⫺1)
䉴
䉴
In the presence of a source of NO, ⭈OH will
therefore be obtained again directly via the oxidation
reaction induced by the HO⭈2 radical, and indirectly via
the photolysis of the ozone produced.
The reactions involving formaldehyde also produce
hydroperoxide radicals, which are among the radical
ENCYCLOPAEDIA OF HYDROCARBONS
ATMOSPHERIC CHEMISTRY
species fundamental in photochemical smog processes
alongside hydroxyl radicals and alkyl peroxide
radicals (formed, as seen above, from organic
compounds).
There is an interconversion between hydroxyl
radicals and hydroperoxide radicals, and the key
reaction in this process is as follows:
[58]
HO⭈⫹NO⫺
⭈OH ⫹NO2
2
䉴
Photochemical pollution is thus caused by a
sequence of interdependent reactions (in some cases
true chain reactions) which give rise to a process
which feeds itself; this explains why acute episodes of
photochemical smog often last for several consecutive
days, increasing in intensity.
Atmospheric secondary particulate matter
The formation of aerosols in the atmosphere has a
significant impact on visibility, the climate and the
chemical processes which occur in the atmosphere; it
is also of special interest since the finest fraction (with
an aerodynamic diameter less than 2.5 mm) may
penetrate the alveoli of the lungs and thus has a direct
impact on human health.
Aerosols in the troposphere may be emitted
directly (primary particulate matter) or be formed
from chemical processes (secondary particulate
matter). The sources are both natural and man made,
and the composition of this material therefore varies
Fig. 3. Negative feedback
cycle between UV radiation
and particulate matter.
considerably. The dimensions of the particles vary
significantly for both primary and secondary
particulate matter, with aerodynamic diameters
ranging from 2 nm to over 10 mm.
Secondary aerosols are generated by the gasparticle conversion which follows the formation,
through oxidative processes, of products characterized
by particularly low volatility or high solubility. Since
these oxidative processes are often of a photochemical
nature, the resulting aerosols may be counted among
the secondary photochemical pollutants.
Secondary aerosol may be generated either by
condensation onto existing aerosol or by nucleation to
form new particles or suspended droplets (Seinfeld,
1986; Clement and Ford, 1996). The most important
molecule which may give rise to the nucleation of new
particles is H2SO4.
Fig. 3 summarizes the processes which occur in
the presence of primary and secondary pollutants
and UV radiation: the interactions between the latter
and the particles have been the object of recent
studies (Kikas et al., 2001). These have shown that
the rate of photolysis of some pollutants (O3, NO2)
depends on the presence of particles which absorb
UV radiation (especially elementary carbon and
organic aerosols; Jacobson, 1999) or scatter it (the
optical properties of the particles vary with size and
are strictly linked to their chemical composition). A
study by Jacobson (1998) has shown that levels of
UV radiation
scattering gases
scattering particulate matter
absorbing gases
absorbing particulate matter
(NO3, HONO, PAN, RO2ⴢ, etc.)
O3
NO2
HOⴢ
SO2
HO⫺, aqueous
oxidation
H2SO4
NO
HNO3
VOC
HO
O3, NO3, HOⴢ
semi-volatile organics
NH3
particulate
sulfate
VOLUME III / NEW DEVELOPMENTS: ENERGY, TRANSPORT, SUSTAINABILITY
particulate
nitrate
secondary
organic aerosol
923
ENVIRONMENTAL TECHNOLOGIES
ozone in Los Angeles have fallen by 5-8% due to
aerosols. Wendisch and colleagues (1996) have
measured the vertical profile of particulate matter to
calculate the vertical rate profile of the NO2
photolysis reaction.
With regard to the composition of secondary
particulate matter, studies have been undertaken
exploiting the different optical properties (light
scattering) at different relative humidities of the
species of interest: H2SO4 and (NH4)2SO4.
These studies have shown that ammonium sulphate
is often observed in real samples, and there is
therefore sufficient ammonia in ambient air to
completely neutralize sulphuric acid (Weiss et al.,
1977). However, sulphate may also be present in other
forms, including ammonium disulphate
(NH4)3H(SO4)2; the formation of mixed salts with
nitrate is unproven, but seems probable under normal
atmospheric conditions.
The equilibria characterizing mixtures of solids
and solutions containing NH4NO3 or mixtures of NH⫹
4,
2⫺ in equilibrium with NH and HNO at
NO⫺
and
SO
3
4
3
3
the concentrations typical of the atmosphere have been
studied and are reported below:
[59]
⫺
NH3(g) ⫹H2SO4(g)⫺
⫺NH⫹
4 ⫹HSO4
䉳
䉴
[61]
2⫺
2NH3(g) ⫹H2SO4(g)⫺
⫺2 NH⫹
4 ⫹SO4
⫺
NH3(g) ⫹HNO3(g)⫺
⫺NH⫹
4 ⫹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)
[60]
䉳
䉳
䉴
䉳
䉳
[66]
䉴
䉳
䉳
䉴
䉴
䉴
䉴
5NH3(g)⫹3 HNO3(g)⫹
⫹H2SO4(g)⫺
⫺(NH4)2SO⭈3NH
4NO3(s)
4
䉳
䉴
500
dC (equivalent / m3) / dDae (mm)
Fig. 4. Distribution
of concentration
of ions with respect
to equivalent diameter.
Examining the composition of particulate matter
⫺
indicates the prevalence of NH⫹
4 and NO3 in the
largest size fraction (diameter between 0.1 and 1 mm);
this is due to the Kelvin effect which involves the
creation of a greater vapour pressure of the volatile
species NH3 and HNO3 on strongly curved surfaces.
This leads to the volatilization of NH3 and HNO3 from
the smallest particles and condensation onto the
largest; this phenomenon does not occur for sulphate,
which has a low volatility.
Fig. 4 shows the distribution of the various species
present in particulate matter as a function of the
equivalent diameter of the particles themselves. Other
components of particulate matter are metals and
organic compounds.
Most Organic Carbon (OC) is found in the fine
fraction of particulate matter. OC derives mainly from
the oxidation of combustion products, such as VOCs,
and their subsequent condensation, dissolution into the
aqueous phase, adsorption (especially onto particles of
elementary carbon, EC), or absorption (Seigneur,
2001). The OC found in the particulate matter emitted
by motor vehicles contains more than 100 different
compounds, including alkanes, benzaldehydes, and
Polycyclic Aromatic Hydrocarbons (PAHs),
particularly dangerous to human health (Rogge et al.,
1993).
The organic compounds of greatest medical
interest are the polycyclic aromatic hydrocarbons:
these form from hydrocarbons with a low molecular
mass by pyrosynthesis at temperatures over 500°C.
The result is the formation of several aromatic rings
condensed into very stable structures which in the
atmosphere, as an effect of solar radiation, may be
transformed into more dangerous compounds such as
nitro-PAHs by reaction with nitric acid and oxidized
PAHs by reaction with ozone. The most frequently
mentioned PAH is benzo(a)pyrene, which
400
300
200
NH⫹
4
SO42⫺
NO⫺
3
Na⫹
Cl⫺
H⫹
100
0 ⫺2
10
10⫺1
1
10
equivalent diameter Dae (mm)
924
ENCYCLOPAEDIA OF HYDROCARBONS
ATMOSPHERIC CHEMISTRY
becomes highly carcinogenic through metabolic
activation. High concentrations of PAHs are present in
the soot generated by the combustion of biomasses
and coal, and by the exhausts of diesel and gasoline
vehicles.
10.1.3 Atmospheric chemistry
in the aqueous phase
Basics
The total volume of water in the atmosphere is
estimated to be about 1.3⭈1013 m3; this water is present
in different forms: aerosols, clouds, fog and rain.
At the northernmost latitudes, over 30% of the
lower troposphere is occupied by cloud bodies. The
water 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 the
type of cloud; in general, their diameter is greater than
10 mm. Fog has a lower water content, about 0.1 g
m⫺3, and smaller droplets.
Aqueous aerosols in the atmosphere consist of
particles with a broad range of diameters, on the basis
of which it is classified as a fine or coarse fraction. A
fine fraction contains free water in various forms:
dilute aqueous solutions, supersaturated solutions and
fine films on insoluble particles.
The volume of the drops present in fog and clouds
is the liquid medium on which the absorption of
reactive trace gases occurs; the majority of these are
highly soluble in water. The concentrations of some
reactive species may therefore be higher in water than
in the surrounding air. This, in conjunction with the
high reaction rates of some species in the aqueous
phase, leads to the conclusion that the droplets
contained in fog and clouds may be extremely efficient
reactors for the oxidation of SO2, NO and NO2.
An important part of atmospheric chemistry takes
place on suspended particles or droplets. The reactions
which occur on the surface or inside these particles are
described as heterogeneous: they take place on the
interface between two phases, i.e. gas-liquid and
gas-solid. Those which take place internally rather
than on the surface are not heterogeneous in the strict
sense of the word, but can be regarded as such
considering the volume of air containing the particles.
The water content in the atmosphere can be
expressed in terms of grams (or of cm3) of water per
m3 of air or as the adimensional fraction of volume L
(for example, m3 of water per m3 of air). Values of L in
different forms of water condensation in the
atmosphere are:
• Clouds, L⫽10⫺7-10⫺6
• Fog, L⫽5⭈10⫺8-5⭈10⫺7
• Aerosols, L⫽10⫺1-10⫺10
VOLUME III / NEW DEVELOPMENTS: ENERGY, TRANSPORT, SUSTAINABILITY
The penetration of atmospheric gases into the
suspended 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 by
Henry’s law, which states that at equilibrium the partial
pressure of a gas on a solution containing the same gas
is proportional to its concentration in solution.
The absorption of a generic species, A, into water
can be represented by the following, entirely
equivalent, equations:
[67] A(g) ⫹H2O⫺
⫺A⭈H2O
䉳
[68]
䉴
A(g)⫺
⫺A(aq)
䉳
䉴
The equilibrium between A in the gaseous form
and the same species in solution can be expressed by
Henry’s constant KH
[69]
KH ⫽[A(aq)]ⲐpA
where pA is the partial pressure of A in the gas phase
and [A(aq)] is the concentration of A in the aqueous
phase at equilibrium. The measurement units for
Henry’s constant are [mol l⫺1 bar⫺1]. Typical values of
Henry’s constant for the main atmospheric gases are
given 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 gases
subsequently react with the water. Henry’s constant
only considers solubility (physical process) and not
hydrolysis reactions (chemical process).
The effect of equilibria in solution is to increase
the quantity of gas passing from the gas phase to the
liquid phase in the atmosphere compared to levels
predicted based on Henry’s law alone.
Given the rapid development of acid-base
equilibria, Henry’s law is often extended by defining a
Henry’s pseudoconstant KH* which takes into account
Table 3. Values of Henry’s constant
for the main atmospheric gases
Gas
KH at 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
1.24
Sulphur dioxide
Ammonia
62
925
ENVIRONMENTAL TECHNOLOGIES
all dissolved species. For example, in the case of SO2
we can define:
[70]
[73]
[74]
[S(IV)] [SO2(aq)]⫹[HSO⫺
3]
11144 ⫽ 1211111124
4⫽
KH*
S(IV)⫽ p
pSO2
SO2
冢
•
冣
K1
K1K2
⫽KHSO2 1⫹ 11
⫹ 112
[H ⫹] [H ⫹]2
where K1 and K2 are the first and second dissociation
constants of the sulphurous acid. It can be seen from
this 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 found
completely dissolved in the tiny water droplets which
form clouds.
The distribution of a generic species, A, between
the gas and aqueous phases in a cloud can be
expressed in terms of the ratio of concentrations of A
in the two phases per unit volume of air:
[71]
moles of A in solution111
per litre
of air 4
1111121121
11124
⫽
moles of A in air per litre of air
H ⭈p ⭈L
K1112
A
⫽1
⫽KH ⭈R⭈T ⭈L
pA ⲐR ⭈T
The constant KH is indicated as KH* if the species
participates in dissociation equilibria in solution.
If KH⭈R⭈T⭈LⰆ1 or KHⰆ1/RTL, species A is present
predominantly in the gas phase, whilst the contrary is
true if KHⰇ1/RTL. Therefore, if L⫽10⫺6,
1/RTL⯝4⭈10⫺4 mol l⫺1 bar⫺1: if Henry’s constant for a
species is less than 4⭈10⫺4 mol l⫺1 bar⫺1 it will be
present mainly in the gas phase.
If the pH is 4 and L is 10⫺6, the value of
H* Ⰶ1/RTL, and SO is therefore
KHS(IV)⯝102 gives KS(IV)
2
present mainly as a gas inside clouds. By contrast, for
a species such as HNO3, KH*⫽1010 and therefore at
equilibrium this species will be present almost entirely
in solution. In acidic environments, H2O2 and NH3 are
also found mainly in the aqueous phase.
The equilibria for the most important species in
atmospheric chemistry in the aqueous phase can be
described as follows:
• carbon dioxide
[72]
CO2⫹H2O⫺
⫺H2CO3(aq)
䉳
䉴
⫺
H2CO3(aq)⫺
⫺H⫹
(aq) ⫹HCO3 (aq)
2⫺
⫺ ⫹
HCO⫺
3 (aq) ⫺H (aq) ⫹CO3(aq)
䉳
䉳
䉴
䉴
sulfur dioxide
[76]
SO2(g) ⫹H2O⫺
⫺H2SO3(aq)
⫺
H2SO3(aq)⫺
⫺H⫹
(aq) ⫹HSO 3(aq)
[77]
2⫺
⫺ ⫹
HSO⫺
3(aq) ⫺H (aq) ⫹SO3(aq)
[75]
•
䉳
䉳
䉳
䉴
䉴
䉴
ammonia
[78]
[79]
NH3 ⫹H2O⫺
⫺NH4OH(aq)
⫺
NH4OH(aq)⫺
⫺NH⫹
4(aq) ⫹OH (aq)
䉳
䉳
䉴
䉴
Henry’s constant KH and the equilibrium constants
for reactions in the aqueous phase (K⬘ and K⬙) are
provided in Table 4.
Both in the case of sulphur dioxide and carbon
dioxide, the second dissociation constant is much
smaller than the first dissociation constant, and can
therefore be disregarded at low pH values.
The pH of a water droplet in equilibrium with
atmospheric CO2 can be calculated by combining the
two equilibrium constants referring to the solubility
and the first dissociation:
[80]
H
⫹
[HCO⫺
3 ] ⫽K ⭈K⬘⭈pCO2 Ⲑ[H ]
If the only source of hydrogen ions is the
⫹
dissociation of carbon dioxide, [HCO⫺
3 ] = [ H ] and
therefore:
[81]
⫹
H
1/2
[HCO⫺
3 ] ⫽[H ] ⫽(K ⭈K⬘⭈pCO2)
Assuming a partial pressure of CO2 of 340 ppm,
this gives a pH of 5.6. This is the pH of rainfall in
remote areas; the presence of traces of other
compounds may affect acidity: SO2 present at a level
of 5⭈10⫺9 bar gives a pH for the solution of 4.6. In
fact, for a solution in equilibrium with SO2, it is
possible to write an equation similar to that for carbon
dioxide:
[82]
H
1/2
[HSO⫺
3 ]⫽(K ⭈K⬘⭈pSO2)
It therefore seems obvious that even low
concentrations of SO2 have a profound effect on pH
although the concentration of CO2 in the atmosphere is
far 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
KH at 288K (mol l⫺1 bar⫺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
Gas
926
K⬙(mol l⫺1)
ENCYCLOPAEDIA OF HYDROCARBONS
ATMOSPHERIC CHEMISTRY
solubility and dissociation constant. These properties
give this species a greater acidifying power which may
even be intensified by concomitant S(IV)⫺ S(VI)
sulphur oxidizing reactions.
When SO2 dissolves in an aqueous solution, the
⫺
2⫺
three resulting species, SO⭈H
2 2O, HSO3 and SO3 ,
with S(IV), may be oxidized by various species:
oxygen (catalyzed by iron and manganese), ozone (the
dominant mechanism when pH⬎5) and hydrogen
peroxide (the dominant mechanism when pH⬍5).
Fig. 5 shows schematically the distribution
equilibrium and the subsequent oxidation equilibria of
sulphur dioxide.
The main oxidation mechanisms are:
• ozone
䉴
Sulphur oxidation
The oxidation of SO2 in the gas phase is mainly
induced by the ⭈OH radical:
[83]
⭈OH⫹SO2 (⫹M)⫺ HOSO2 (⫹M)
[84]
⫺ ⫺ H2SO4
HOSO⭈⫺
2
䉴
䉴
䉴
䉴
The ⭈OH radical forms in the atmosphere in the
presence of mixtures containing non-methane
hydrocarbons and NOx due to the effect of solar
radiation and gives rise, through the formation of the
aforementioned radical species, to the formation of
sulphuric acid. There are various hypotheses regarding
the mechanism by which the radical species formed
gives rise to H2SO4 (Benson, 1978; Calvert et al.,
1978; Davis et al., 1979); this discussion consequently
makes use of the simplified equation above.
Oxidation in the gas phase induced by the ⭈OH
radical is characterized by formation rates greater than
1% an hour; studies conducted both in the field and in
the laboratory have recorded values for the conversion
rate higher than those predictable on the basis of gas
phase chemistry alone. This leads to the conclusion
that reactions in the aqueous phase may be important
sources of sulphate and in some cases may even be
more important than those in the gas phase.
[85]
•
⫺
2⫺
HSO⫺
3(aq) ⫹OH ⫹O3⫺ SO4 ⫹H2O ⫹O2
䉴
hydrogen peroxide
[86]
⫺
⫺
HSO⫺
3(aq) ⫹H2O2(aq) ⫺SO2OOH ⫹H2O
[87]
SO2OOH⫺ ⫹H⫹⫺ H2SO4
•
䉳
䉴
䉴
oxygen
[88]
Fe, Mn
2⫺ ⫹H⫹
HSO⫺
⫺ SO4(aq)
3(aq) ⫹1Ⲑ2O2(aq)⫺
(aq)
䉴
Ozone reacts very slowly with SO2 in the gas
phase, whereas the reaction is fast in the aqueous
phase. The most plausible mechanism is the ionic
mechanism shown above and proposed by Maahs
(1983). Radical mechanisms have also been
proposed whose contribution to the oxidation of
S(IV) is not quantifiable with any degree of
certainty. These may take place via a radical
intermediate such as ⭈OH (Hoignè and Bader, 1975;
Penkett et al., 1979):
SO2(g)
Fig. 5. Sulphur dioxide
distribution and oxidation
equilibria.
(1)
transport to droplet surface
(2)
transport across air-water interface
SO2(interface)
SO2⭈H2O
HSO⫺
3
(3)
establishment of S(IV) equilibria
SO2⫺
3
(4)
S(VI)
transport into bulk phase
S(IV)
(5)
VOLUME III / NEW DEVELOPMENTS: ENERGY, TRANSPORT, SUSTAINABILITY
oxidation
927
ENVIRONMENTAL TECHNOLOGIES
[89]
[90]
[91]
⫺
HSO⫺
3 ⫹⭈OH⫺ H2O ⫹SO3
⫺
SO⫺
3 ⫹O2⫺ SO5
⫺
⫺
2SO5 ⫺ SO4 ⫹SO2⫺
4 ⫹O2
䉴
䉴
䉴
Oxidation induced by H2O2 is relatively
independent of pH and this is due to the fact that the
rate constant of the reaction and the solubility of S(IV)
show opposing trends with respect to pH. Other
species are affected in different ways by variations in
pH: in the case of ozone the reaction rate increases by
one order of magnitude when passing from pH 1 to pH
3, whereas the reaction rate of oxygen catalyzed by
iron increases by two orders of magnitude. The
dependence of rate and the solubility of S(IV) on pH
means that the reaction with ozone is only important
when pH⬎5.
Research has concentrated particularly on reactions
in the aqueous phase catalyzed by transition metals:
the most important catalysts are iron (in the form
Fe3⫹) and manganese (in the form Mn2⫹). Some
studies have been carried out on the possible catalytic
action of other metals such as Cu2⫹ and Co2⫹;
however, the reactions catalyzed by the latter are
characterized by very low reaction rates under typical
ambient conditions.
The rate of the reaction catalyzed by Fe(III)
depends on various parameters, such as pH, ionic
force and temperature, and is affected by the presence
of some anions (such as SO42⫺) and cations (Mn2⫹) in
solution.
The mechanisms of these reactions have long been
debated. A radical mechanism has been suggested
(Hoffmann and Boyce, 1983; Hoffmann and Jacob,
1984):
• initiation
(n⫺1)⫹⫹SO⭈⫺
[92] Mn⫹⫹SO2⫺
3 ⫺M
3
䉴
•
propagation
[93]
⫺
SO⭈⫺
3 ⫹O2⫺ SO⭈5
[94]
2⫺
2⫺
⫺
SO⭈⫺
5 ⫹SO3 ⫺ SO5 ⫹SO⭈3
•
䉴
䉴
metal-sulphite complex followed by bonding with
oxygen:
2⫺
⫺
Mn2⫹⫹2SO2⫺
3 ⫺Mn(SO3)2
2⫺
⫺
[100] Mn(SO3)2⫺
2 ⫹O2 ⫺Mn(SO3)2O2
2⫹
2⫺
⫺
[101] Mn(SO3)2O2⫺
2 ⫺Mn ⫹2SO4
[99]
䉳
䉴
䉳
䉳
䉴
䉴
Mechanisms that are even partially photochemical
have been suggested, based on the observation (Lunak
and Veprek-Siska, 1976) that the oxidation of S(IV) in
homogeneous aqueous systems for wavelengths above
300 nm does not occur unless Fe3⫹ is present.
Photooxidation has been attributed to the absorption of
light by a Fe3⫹- S(IV) complex.
In the case of iron, this is present in solution in the
soluble and solid form:
[102] [Fe]tot ⫽[Fe]sol ⫹[Fe(OH)3]
Soluble iron is found in aqueous solution in
various ionic forms:
[103] [Fe]sol ⫽[Fe3⫹] ⫹[FeOH2⫹] ⫹[Fe(OH)⫹2 ] ⫹
⫹[Fe2(OH)24⫹]
The equilibrium between iron in the soluble form
and solid iron is as follows:
[104] Fe(OH)3 ⫹3H⫹⫺
⫺Fe3⫹ ⫹3H2O
䉳
䉴
As a result, the relative quantities of the different
forms of iron in aqueous solution are strongly
influenced by pH. With pH values of 4.5,
[Fe3⫹]⯝3⭈10⫺11; above this pH value the soluble form
of iron decreases significantly.
Another catalyst for these oxidation reactions is
manganese: attempts have been made to determine
which of the two metals is the main cause of SO2
oxidation, but to date their relative importance in the
reaction has not been clarified.
It is likely that the two catalysts act in synergy: an
increase in the reaction rate has been observed in the
presence of both ions, higher than would be expected
on the basis of the sum of their respective catalysis
oxidation
[95]
2⫺
2⫺
SO2⫺
5 ⫹SO3 ⫺ 2SO4
䉴
• termination
2⫺
[96] 2SO⭈⫺
3 ⫺ S2O6
⫺
2⫺
[97] SO⭈⫺
3 ⫹SO⭈5 ⫺ S2O6 ⫹O2
⫺
⫺
[98] SO5 ⫹SO5 ⫺ S2O82⫺ ⫹O2
䉴
Table 5. Concentrations (expressed in molarity, M)
of iron and manganese in various aqueous matrices
present in the atmosphere
䉴
䉴
This hypothesis was later abandoned since it
clashed with the kinetic order of reaction of the
reagents and was replaced by two alternative
mechanisms: ionic and photochemical.
The ionic mechanism, initially proposed by Bassett
and Parker (1951), involves first the formation of a
928
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
ENCYCLOPAEDIA OF HYDROCARBONS
ATMOSPHERIC CHEMISTRY
rates: the removal of S(IV) is about 3–10 times faster
if there is synergy between the two species (Martin,
1984). Table 5 quantifies the presence of these two
species in the various forms of water condensation in
the atmosphere.
Temperature has the opposite effect on these
catalyzed reactions compared to other oxidation
reactions: the formation of sulphate via these
mechanisms decreases as temperature decreases. In
fact, the presence of the catalysts does not vary with
temperature, as is the case for the other oxidizing
species found in the gaseous form, whilst the high
activation energies are affected by decreases in
temperature.
Oxidation by Mn2⫹ is inversely influenced by
pH with respect to Fe3⫹: whereas in the case of
iron the reaction rate decreases as pH decreases in
the range 0-4, the opposite occurs for manganese
in the range 0-3.
Fig. 6 shows the effect of pH on the various types
of oxidation. It can be observed that the dominant
mechanism for the formation of sulphate at pH lower
than 4-5 is that induced by hydrogen peroxide. In
contrast, when pH⭓5 oxidation by ozone is 10 times
faster than that by H2O2. Oxidation catalyzed by
metals is important when the pH is high. Assuming a
water content of 1 g m⫺3 within a cloud, the rate of
oxidation induced by H2O2 exceeds 100% an hour,
whereas the rates of oxidation reactions catalyzed by
iron and manganese are lower than 1% an hour when
pH ⬍4.5.
104
O3, 50 ppb
103
H2O2, 1 ppb
rate of S(IV) oxidation (%/h)
102
Fe, 3ⴢ10⫺7mol/l
101
Mn, 3ⴢ10⫺8mol/l
The oxidizing power of nitrogen oxides with
respect to S(IV) differs depending on the species
under consideration. NO and HNO3 induce reactions
which are too slow to contribute effectively to
oxidation; HONO, on the other hand, has extremely
low atmospheric concentrations (1-8 ppb) and, despite
its high Henry’s constant (KH⫽49 mol l⫺1 bar⫺1), does
not reach sufficient concentrations in solution to
represent an important oxidizing agent.
NO2, whose Henry’s coefficient is KH⫽1⭈10⫺2 mol
l⫺1 bar⫺1, by contrast, is a relatively insoluble gas;
however, it has been shown (Schwartz, 1984) that the
2⫺
reaction rates of NO2 with HSO⫺
3 and SO3 are
sufficiently high to make its oxidizing action important.
Additionally, it has been shown that the reaction:
⫹
⫺
2⫺
[105] 2NO2⫹HSO⫺
3 ⫺ 3H3O ⫹2NO2 ⫹SO4
䉴
may be sufficiently fast at high pH levels (Lee and
Schwartz, 1983; Lee, 1984).
Oxidation of nitrogen
In addition to oxidation from S(IV) to S(VI),
another possible oxidation process leading to the
acidification of the water present in clouds or rain is
that induced by NO2 and NO and giving rise to nitric
and nitrous acid:
⫺
⫺
[106] 2NO2(g) ⫹H2O(l)⫺ 2H⫹
(aq) ⫹NO 3(aq) ⫹NO 2(aq)
䉴
⫺
[107] NO(g) ⫹NO2(g) ⫹H2O(l)⫺ 2H⫹
(aq) ⫹2NO 2(aq)
䉴
However, these reactions do not contribute
significantly to acidification at the concentrations of
nitrogen oxides normally present in the atmosphere:
the first of the two reactions (Schwartz, 1984) is very
slow, both due to the low solubility of NO2 and to the
second order dependence of the reaction rate on the
concentration of NO2.
The formation mechanism for nitric acid in air
remains primarily the same as the gas phase induced
by the free radical ⭈OH:
[108] ⭈OH ⫹NO2 (⫹M)⫺ HONO2 (⫹M)
䉴
1
C, 10⫺2g/l
10⫺1
HNO2, 1 ppb
NO2, 2 ppb
10⫺2
10⫺3
10⫺4
10⫺5
Analysing the composition of precipitation and the
surrounding air masses nevertheless shows that
oxidation in the aqueous phase also produces
significant quantities of nitric acid in the atmosphere
(Lazrus et al., 1983; Misra et al., 1985).
A mixed gas-liquid mechanism for the production
of nitric acid in droplets has been proposed by Heikes
and Thompson (1983):
[109] O3 ⫹NO2⫺ O2 ⫹NO3
䉴
0
1
2
3
4
5
6
pH
Fig. 6. Effect of pH on various sulphur(IV)
oxidation reactions.
VOLUME III / NEW DEVELOPMENTS: ENERGY, TRANSPORT, SUSTAINABILITY
[110] NO3 ⫹NO2 (⫹M)⫺ N2O5 (⫹M)
䉴
N2O5 generated in this way in the gas phase may be
incorporated into droplets or onto the surface of
aqueous aerosol, and hydrolyze to HNO3:
929
ENVIRONMENTAL TECHNOLOGIES
⫺
[111] N2O5 ⫹H2O⫺ 2H⫹
(aq) ⫹2NO 3(aq)
䉴
However, a definitive evaluation of the
contribution of this mechanism to the concentration of
HNO3 in the liquid phase of the atmosphere is not yet
possible.
Special attention has recently been devoted to
reactions in the aqueous phase involving species
generated by photochemical mechanisms: inside cloud
bodies there may be sufficient light for the production
of hydroxyl and hydroperoxide radicals.
Hydroxyl radicals react with nitrogen oxides to
form nitrous and nitric acid:
⫺
[112] NO⫺
2(aq) ⫹⭈OH(aq)⫺ NO2(aq) ⫹OH (aq)
䉴
[113] NO(aq) ⫹⭈OH(aq)⫺ HNO2(aq)
䉴
[114] NO2(aq) ⫹⭈OH(aq)⫺ HNO3(aq)
䉴
Other potential reactions leading to the formation
of nitric acid are those suggested by Heikes and
Thompson (1983) and Platt et al. (1984): the nitrate
radical absorbed in solution is hydrolyzed to form
nitric acid. This may occur starting from the radical
itself with the extraction of a hydrogen atom from an
organic compound:
[115] NO⭈⫹RH⫺
R⭈⫹HNO3
3
䉴
The reaction with H2O2, however, does not appear
to be important:
⫹
[116] HONO ⫹H2O2⫺ NO⫺
3 ⫹H3O
䉴
since it has been demonstrated (Lee, 1984) that this
reaction is too slow.
In the absence of clouds and in the presence of
solar radiation, HNO3 is generated in the gas phase by
the reaction of NO2 with ⭈OH radicals at a rate of
about 20-30% per hour, whilst H2SO4 is generated by
this pathway far less effectively. By contrast, in the
aqueous phase oxidation reactions involving H2O2, O3
and metal catalysts are able to produce H2SO4 in
solution at rates of up to 100% an hour.
The chemico-physical properties of these two acid
species are fundamentally different: nitric acid is far
more volatile and tends to remain in the atmosphere in
the gaseous form, whilst sulphuric acid has a relatively
low vapour pressure (⬍10⫺7 bar) and therefore tends
to be present inside particles. Both may react with
basic species. These reactions essentially occur with
the ammonia present to form ammonium nitrate and
ammonium sulphate:
[117] HNO3 ⫹NH3⫺
⫺NH4NO3
䉳
䉴
[118] H2SO4 ⫹2NH3⫺
⫺(NH4)2SO4
䉳
䉴
These two species are the main cause of
diminishing visibility in the case of photochemical
930
smog. Ammonium nitrate is found in the solid form if
the temperature is below that of deliquescence; if
relative humidity is high, it is present in the aqueous
phase. However, nitric acid may return to the gas phase
even after forming ammonium salts.
Importance of aqueous phase equilibria
in the formation of acid rain
Formed in the atmosphere by the mechanisms
described above in the aqueous and gas phases, the
acids under examination may be deposited on the
Earth’s surface by two main mechanisms: dry
deposition and wet deposition. The distinction is made
based on the phase in which the pollutant is found
when it hits the surface: dry deposition involves
gaseous pollutants or small particles, whilst wet
deposition involves pollutants present in the droplets
inside fog, clouds and rain.
It should be specified that this classification only
considers the transportation mechanism, and not the
nature of the surface (whether or not it is wet or if it
presents a liquid film). Given the variable nature of
precipitation it is difficult to make quantitative
calculations of the extent of wet deposition of a
pollutant species.
The deposition rate of a pollutant is given
approximately by the product l⭈C, where C is the
concentration of the pollutant and l is the scavenging
coefficient proportional to the intensity of the
precipitation itself.
Dry deposition, in contrast, is characterized by the
deposition rate Vg, defined as:
[119] Vg ⫽FⲐ[S]
where F is the flux of species S towards the surface
and [S] is the concentration of the same species at a
reference altitude h.
As far as wet deposition by precipitation is
concerned, the sulphate and nitrate deposited on
exposed surfaces may be the result not only of
oxidation reactions in the aqueous phase, but also of
the inclusion of aerosol particles containing these
species inside the cloud or of the scavenging of
aerosol particles beneath the cloud itself. If these
inclusion and scavenging mechanisms predominate,
acidification will be determined by the parameters
governing the formation of nitrates and sulphates in
the gas phase: solar radiation, the concentration of
NOx and hydrocarbons.
If formation inside the droplets by oxidation of
SO2 and NO2 prevails, however, the parameters
determining acidity will be the concentrations of
oxidants such as H2O2 and O3.
Comparing the composition of interstitial aerosols (a
fraction of the atmospheric aerosol which, in the
ENCYCLOPAEDIA OF HYDROCARBONS
ATMOSPHERIC CHEMISTRY
presence of a cloud, remains the same and is not
removed by the water droplets present inside the cloud)
and that of clouds provides information on the source of
acidity inside the clouds themselves. Daum and
colleagues (1983) have shown that for interstitial
aerosols the ratio [H⫹]/[NH⫹
4 ] is less than 1, whereas the
aerosol contained in clouds has a ratio greater than 1.
Similar conclusions have been reached by other
research groups (Lazrus et al., 1983; Harrison and Pio,
1983): the higher acidity of the aerosols in cloud bodies
with respect to that of the surrounding atmosphere
agrees well with the formation mechanisms of acidic
species in the aqueous phase described above.
It can be concluded that the formation of sulphuric
acid inside cloud aerosols provides the largest
contribution to the formation of sulphate in acid rain.
10.1.4 Depletion of stratospheric
ozone
The stratosphere is that region of the Earth’s atmosphere
which stretches from about 15 to 50 km above the
troposphere. The temperature is almost constant in the
lowest layer, whilst it tends to increase gradually in the
upper half. The temperature pattern in the stratosphere
is regulated by variations in the concentration of ozone
inside this layer: by absorbing solar radiation in the
ultraviolet band ozone molecules convert solar energy
into kinetic energy, helping to heat the stratosphere.
The average global concentration of stratospheric
ozone varies with altitude; at an altitude of 15 km it is
in the order of 0.5 parts per million (ppm) by volume,
at around 35 km it increases up to 8 ppm and then
decreases to 3 ppm in the high stratosphere (45 km).
The thickness of the ozone layer above a certain
geographic area or its average global value can be
conventionally expressed in Dobson Units (DU), in
other words with a height of 0.01 mm; the ozone would
be at this thickness if it were the only component of the
atmosphere and if it were at a pressure of 1 bar and at a
temperature of 0°C. On average the height of the ozone
layer may vary with latitude between 250 and 400 DU;
its average global value is about 300 DU.
The concentration of ozone in the various regions
of the stratosphere is the result of a dynamic formation
and destruction process. Ozone forms at an altitude of
about 30 km, where ultraviolet solar radiation with a
wavelength shorter than 242 nm slowly dissociates the
oxygen molecules into atomic oxygen:
[121] O ⫹O2 ⫹M⫺ O3 ⫹M
䉴
In turn, the ozone molecules absorb the high
energy photons present in solar radiation, with
wavelengths ranging from 240 to 320 nm. A direct
consequence of this absorption is the dissociation of
ozone into an oxygen molecule and an excited oxygen
atom, O(1D), according to the reaction:
[122] O3 ⫹hv⫺ O2 ⫹O
䉴
This absorption process means that a significant
portion of ultraviolet solar radiation does not reach the
Earth’s surface: this has made the development of life
possible on our planet. The photolysis of ozone is not a
genuine destruction mechanism because virtually all
the oxygen atoms produced by this reaction combine
rapidly with oxygen molecules to form ozone again.
This mechanism entails the conversion of solar energy
into thermal energy, especially in the upper part of the
stratosphere. The presence of ozone is thus the cause
of the temperature inversion characterizing the upper
belt of the stratosphere.
The current distribution of stratospheric ozone is
mostly determined by transportation processes. Ozone
is produced mainly at the tropics at an altitude of
between 25 and 35 km but, as a result of the
movements of air masses, its highest concentrations
are found near the poles at an altitude of about 15 km.
Various destruction mechanisms contribute to
balancing out ozone formation processes (reactions
[120] and [121]). An example is the reaction of ozone
with oxygen atoms to form molecular oxygen:
[123] O3 ⫹O⫺ O2 ⫹O2
䉴
The above reactions are known as Chapman
reactions and have formed the basis for the study of
stratospheric ozone. Chapman’s scheme only
estimates the loss of ozone through natural
destruction and does not consider the transportation
of ozone onto the Earth’s surface which contributes
a further 0.5%.
About 10% of destruction processes can be
attributed to catalytic cycles involving species which
contain hydrogen: free hydrogen atoms (H), hydroxyl
(⭈OH) and hydroperoxide (HO⭈)
2 radicals give the
same results as reaction [123].
For example:
⭈OH ⫹O2
O ⫹HO⭈⫺
2
HO⭈⫹O3⫺ HO⭈⫹O
2
2
11111111122
O3 ⫹O⫺ O2 ⫹O2
These species containing hydrogen are generated
by the reaction which normally occurs between water
vapour and methane and the excited oxygen atoms O
(1D) from the photolysis of ozone.
䉴
䉴
䉴
[120] O2 ⫹hv⫺ O ⫹O
䉴
The oxygen atoms rapidly combine with molecular
oxygen to form ozone; in the presence of a third inert
molecule, M, the following reaction may take place:
VOLUME III / NEW DEVELOPMENTS: ENERGY, TRANSPORT, SUSTAINABILITY
931
ENVIRONMENTAL TECHNOLOGIES
A significant contribution, about 70%, to the
process of ozone destruction is supplied by a catalytic
cycle involving NO and NO2, the main reactions of
which are:
O⫹NO2⫺ NO⫹O2
NO⫹O3⫺ NO2⫹O2
11111111122
O3⫹O⫺ O2⫹O2
The main source of NOx in the stratosphere is the
oxidation of the nitrous oxide (N2O) produced by
bacteria in the land and water. Although most nitrous
oxide 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 on
ozone to form nitrogen oxide and begin the cycle:
䉴
Table 6. Halogenated compounds present
in the troposphere
Name
Chemical
formula
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
Compound
䉴
䉴
[124] O (1D) ⫹N2O⫺ NO ⫹NO
䉴
The NOx molecules, in turn, are removed by the
formation of nitric acid (HNO3) in the reaction
⭈OH-NO2 which occurs in the lowest layer of the
stratosphere. Atmospheric currents transport the nitric
acid towards the troposphere, where it is removed by
precipitation; in the same way, the nitrous oxide from
the troposphere is transported upwards and destroyed.
The destruction of ozone may also be catalyzed by
substances other than HOx e NOx, in particular by
chlorine (Cl) and bromine (Br) atoms and their
respective oxides (ClO and BrO).
The year 1930 saw the beginnings of the industrial
production of chlorofluoromethanes and
chlorofluoroethanes, known by the name freon, which
became widely used as liquid refrigerants, solvents
and propellants for aerosol cans. In 1975, two
researchers, F.S. Rowland and M.J. Molina, published
an article stating that in the stratosphere freon could
induce radical chain reactions, with negative effects
the on the natural equilibrium of ozone. An example of
the reactions taking place is as follows:
• initial stage
[125] CF2Cl2 ⫹hv⫺ CF2Cl⭈⫹Cl
䉴
•
propagation stage
[126] O ⫹ClO⭈⫺ Cl⭈⫹O2
䉴
[127] Cl⭈⫹O3⫺ ClO⭈⫹O2
䉴
In the initial stage, ultraviolet light causes the
homolytic cleavage of a C⫺Cl bond of the freon. The
chlorine radical may induce the propagation of a chain
reaction leading to the destruction of ozone molecules.
Again in 1976, a study by the National Academy of
Sciences (NRC, 1976) confirmed Rowland and
Molina’s predictions and in January 1978 the use of
freon in aerosol cans was banned in the United States.
The main chlorine compounds present in the
troposphere (Table 6) are methyl chloride (CH3Cl), a
932
tiny proportion of which is of industrial origin, the
man-made chlorofluoromethanes CFCl3 (CFC-11) and
CF2Cl2 (CFC-12), and carbon tetrachloride (CCl4)
generated by both natural and man-made sources. Less
important man-made sources include trichloroethylene
(CCl2⫽CHCl) and the substances which have
replaced it, methyl chloroform (CH3CCl3) and
trichlorotrifluoroethane (C2F3Cl3), CFC-113.
From the point of view of their effects on the
stratosphere, the key species are the
chlorofluorocarbons CFCl3 (CFC-11) and CF2Cl2
(CFC-12) whose concentrations in the atmosphere
increased by 37% and 31% respectively between 1976
and 1981. It is thought that the residence times of
CFC-11 and CFC-12 are about 60 and 110 years
respectively; measurements indicate a minimum
life-time for CFC-11 of 40 years.
A few years after their commercialization, these
chlorine compounds spread through the troposphere,
and their concentration subsequently began to increase
slowly in the stratosphere as well.
Chlorofluorocarbons (CFC-11, CFC-12, CFC-113) are
highly inert compounds in the troposphere and in the
lower layers of the stratosphere, but when transported
to altitudes of 25-50 km they are decomposed by
ultraviolet radiation with wavelengths shorter than
about 200 nm, with the resulting formation of chlorine
atoms.
The chlorine atoms subsequently participate in the
catalytic cycle of ClOx. This cycle may be interrupted
by the conversion of the highly reactive forms Cl and
ClO into less reactive forms which do not destroy
ozone. The chlorine atoms are deactivated by reaction
with methane to form HCl:
[128] Cl⫹CH4⫺ HCl⫹CH3
䉴
ENCYCLOPAEDIA OF HYDROCARBONS
ATMOSPHERIC CHEMISTRY
which acts as a temporary reservoir for the active
chlorine species in the stratosphere. The chlorine
atoms are regenerated by the reaction between HCl
and ⭈OH radicals:
[129] ⭈OH ⫹HCl⫺ H2O ⫹Cl
䉴
The destruction and regeneration processes of the
active chlorine species, ClOx , may occur several times
before the chlorine is completely removed from the
stratosphere. The most important removal mechanism
is the transportation of HCl from the stratosphere to
the upper troposphere, from which it is removed by the
action of rain, as occurs for the removal of nitric acid
from the NOx cycle. The time scale from the point at
which the chlorofluorocarbons are emitted at ground
level to the point at which their chlorine atoms, in the
form of HCl, are removed from the atmosphere by rain
is in the order of several decades; variations in the
emission fluxes of chlorofluorocarbons thus manifest
themselves in the stratosphere after many years.
Ozone destruction reactions may also be catalyzed
by bromine atoms. It is thought that stratospheric
bromine is 50 times more active than Cl in the
processes leading to the destruction of ozone, and that
it is responsible for 20% of the ozone hole over the
Antarctic and for a greater proportion of that over the
Arctic. The most important transportation vehicle for
bromine is methyl bromide (CH3Br); the photolysis of
this compound leads to the formation of the Br radical,
responsible for the demolition of O3:
[130] CH3Br ⫹hv⫺ ⭈CH3 ⫹⭈Br
In order to replace chlorofluorocarbons, other
substances have been synthesized for use as
refrigerants and aerosol propellants:
hydrochlorofluorocarbons (HCFCs) and
hydrofluorocarbons (HFCs). The presence of at least
one hydrogen atom in each molecule means that these
are oxidized in the troposphere by reactions with ⭈OH
radicals:
[134] ⭈OH ⫹CHClxFy⫺ H2O ⫹CClxFy⫺ products
䉴
䉴
The presence of hydrogen thus allows for a faster
degradation of the substance in the atmosphere, and
therefore a lower environmental impact (HCFCs and
HFCs do not reach the stratosphere).
The Ozone Depleting Potential (ODP) has been
introduced to combat the harmful effects of the
chlorofluorocarbons and the substances which have
replaced them. For any given halocarbon, the ODP is
defined as the ratio of stratospheric ozone destroyed
by the emission of 1 kg of that compound to the ozone
destroyed by the emission of 1 kg of CFC-11.
A substance’s ozone depleting potential thus
provides a measure of the impact (in comparison to
that of CFC-11, considered the standard reference
compound with an ODP of 1.00) of the emission of 1
kg of that substance in terms of the destruction of
stratospheric ozone. As such, the impact of a substance
on the ozone layer is given by its capacity to destroy
ozone and the quantity of total emissions. Table 7
shows the ODP values and life-times in the
atmosphere of the main compounds.
䉴
[131] Br ⫹O3⫺ BrO⭈⫹O2
䉴
[132] BrO⭈⫹O⫺ ⭈Br ⫹O2
䉴
Little is known about methyl bromide’s sources and
cycle, unlike chlorofluorocarbons. This gas is used as
a fumigant for the antiparassitic treatment of land and
agricultural products and as a raw material for
products of chemical synthesis. Together, natural sources
(mainly marine biological activity; Lovelock, 1975) and
man-made sources emit about 100 kt a year into the
environment. Considering the emission fluxes and a
life-time of about 2 years, the average global concentration
of methyl bromide is on the order of 10 ppt.
A particularly interesting aspect of the chemistry
of stratospheric bromine is the possible synergic
interaction between its cycle and that of chlorine,
through the reaction:
[133] BrO ⫹ClO⫺ Br ⫹Cl ⫹O2
Table 7. ODP values and life-times
of the main halogenated compounds present in the
atmosphere (WMO/United Nations Environment
Programme, 1994)
ODP
Life-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
Compound
䉴
Yung and colleagues (1980) suggest that the
interaction between ClOx and BrOx may lead to an
increase in the ozone destruction process in the lower
layers of the stratosphere.
VOLUME III / NEW DEVELOPMENTS: ENERGY, TRANSPORT, SUSTAINABILITY
933
ENVIRONMENTAL TECHNOLOGIES
Other reactions may occur in the heterogeneous
phase on the surface of the particles present in Polar
Stratospheric Clouds (PSCs), consisting of an icy
mixture of nitric acid and water, generated at
temperatures below 195K (⫺78°C); these clouds,
pearly white in colour, can frequently be seen over the
Antarctic. HCl molecules aggregate onto the ice
crystals, becoming part of the crystalline structure and
leading to the formation of a solid phase catalyst. The
catalyst reacts with chlorine nitrate to release chlorine
in the gaseous state; this accumulates in the
stratosphere and participates in the ozone destruction
process:
[135] ClONO2 ⫹HCl⫺ Cl2 ⫹HNO3
䉴
The nitric acid remains bonded to the surface of
the particles and therefore reacts with NO2, which is
no longer able to neutralize ClO. In this way, the level
of active chlorine atoms is kept constant.
The low temperatures reached in polar Antarctic
regions are associated with low pressure conditions
leading to the formation of a polar vortex which mixes
troposphere and stratosphere whilst simultaneously
preventing the entry of external air masses. The same
low temperatures encourage the formation of PSCs in
the Antarctic region, a phenomenon of minor
importance in the Arctic, where temperatures are about
10-15°C higher.
It appears that aerosol particles containing
sulphuric acid may present surface conditions that
accelerate the ozone destruction process. This was
shown clearly following the high emissions into the
stratosphere of sulphur compounds in particulate and
gaseous form by explosive volcanic eruptions (an
example is the eruption in 1991 of Mount Pinatubo in
the Philippines).
The release of large amounts of SO2 leads to the
formation of aerosol through the reaction:
the water vapour present in small quantities in the
stratosphere to form aerosol droplets.
Volcanoes are only one source of the gaseous
sulphur present in the stratosphere; another, of
biogenic origin, releases carbon disulphide. The levels
of carbon disulfide present in the stratosphere are
extremely low, but it may be transferred there across
the tropopause, unlike other sulphur compounds such
as hydrogen sulphide and sulphur dioxide, apparently
too reactive to leave the troposphere. The high
concentrations of particulate matter containing
sulphate (in the form of both SO42⫺ and H2SO4) are
shown in Fig. 7.
The diagram shows that carbonyl sulphide (OCS)
plays an important role in the chemistry of the sulphur
compounds present in the atmosphere. It can be
oxidized to sulphuric acid:
[140] OCS ⫹⭈OH⫺ CO2 ⫹HS
䉴
[141] HS ⫹⭈OH⫺ SO ⫹H2
䉴
[142] SO ⫹⭈OH⫺ SO2 ⫹H
䉴
It has been estimated that over 50% of the carbonyl
sulphide present in the atmosphere is man made and is
emitted by combustion processes and the treatment of
fossil fuels.
Some models suggest that one of the most
significant consequences of ozone depletion will be
the cooling of the lower part of the stratosphere. It is
thought that there is a positive feedback mechanism by
which the loss of ozone cools the air, encouraging the
formation of stratospheric polar clouds which in turn
40
H2SO4
30
[136] ⭈OH ⫹SO2⫺ HSO3
which may considerably decrease the
concentrations of ⭈OH radicals in the upper part of the
atmosphere, with notable consequences. The
subsequent reactions involving HSO3 are not fully
understood, but a potential further step might be:
[137] HSO3 ⫹⭈OH⫺ H2SO4
䉴
altitude (km)
䉴
SO42⫺
SO2
20
tropopause
OCS
10
CS2
similarly
H2S and
(CH3)2S
Alternatively, two other reactions could be
suggested:
[138] HSO3 ⫹O2⫺ HSO5
䉴
[139] HSO3 ⫹O2⫺ HO⭈⫹SO
3
2
䉴
The oxidation of SO2 of volcanic origin in the
atmosphere takes place very slowly. The sulphuric acid
generated in this way may subsequently condense on
934
0
10⫺12
10⫺11
10⫺10
10⫺9
concentration (molar fraction)
Fig. 7. Concentration of the main sulphur compounds
as a function of altitude.
ENCYCLOPAEDIA OF HYDROCARBONS
ATMOSPHERIC CHEMISTRY
contribute to the lowering of ozone levels. It is obvious
that the formation of PSCs is not limited to polar
vortices, but may occur inside stratospheric jet streams
at temperate latitudes. All this suggests that
heterogeneous processes involving aerosols containing
ice particles, sulphuric acid and sulphur compounds
are important components of the chemistry of the
stratosphere.
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Ivo Allegrini
Consiglio Nazionale delle Ricerche
Istituto sull’Inquinamento Atmosferico
Monterotondo, Roma, Italy
ENCYCLOPAEDIA OF HYDROCARBONS