Below are the first 10 and last 10 pages of uncorrected machine-read text (when available) of this chapter, followed by the top 30 algorithmically extracted key phrases from the chapter as a whole.
Intended to provide our own search engines and external engines with highly rich, chapter-representative searchable text on the opening pages of each chapter.
Because it is UNCORRECTED material, please consider the following text as a useful but insufficient proxy for the authoritative book pages.
Do not use for reproduction, copying, pasting, or reading; exclusively for search engines.
OCR for page 155
Appendix A
The Chemistry of
Acid Formation
It is well established that the oxides of sulfur--sulfur
dioxide (SO2) and sulfur trioxide (SO3)--and nitrogen--
nitric oxide (NO) and nitrogen dioxide (NO2)--are
converted (oxidized) in the troposphere to sulfuric acid
(H2SO4) and nitric acid (HNC3), respectively. However,
the most prevalent end products of these reactions--H2SO4,
HNO3, ammonium bisulfate (NH4HSO4), ammonium nitrate
(NH4NO3), etc.--give few clues about which of several
oxidizing pathways are important. Yet for the develop-
ment of scientifically sound, predictive models of acid
deposition that define theoretical source-receptor rela-
tionships r knowledge of the elementary chemical steps
that are involved (among many other quantities such as
deposition rates and transport rates) is required. The
rates of these reactions follow well-defined mathematical
laws (rate expressions) that relate reaction rates to
concentrations of the reactants, temperature, pressure,
and other variables. Considerable progress has been made
in developing an understanding of the nature of this
chemistry, but a large number of uncertainties still
remain in key areas.
As our knowledge of the atmospheric chemistry of
SO2, NO, and NO2 continues to grow, it has become
increasingly clear that many different pathways exist for
generating H2SO4 and HNO3 in the troposphere.
Reactions can occur in the gas phase; in the solution
phase in cloud water and rainwater, for example; and in
reactions on surfaces of solid particles in the
atmosphere. Thus we are concerned with the rates of
exchange of gaseous reactants and their reaction products
between liquids and the surfaces of solids as well as the
rates of interactions among gaseous molecules, aqueous
phase molecules and ions, and species adsorbed on solid
155
OCR for page 156
156
surfaces. For current purposes, it is not necessary to
review and evaluate every detail of these chemical
processes, but we should note the nature of the many
chemical processes, the variety of interactions among the
many species involved, and the current state of knowledge
related to the chemistry of acid deposition. We consider
first the homogeneous gas-phase chemistry that results in
oxidation of SO2, NO, and NO2 in the troposphere (Calvert
and Stockwell 1983, Calvert et al. 1978).
GAS-PHASE REACTIONS LEADING TO
GENERATION OF ACID IN THE TROPOSPHERE
Oxidation by Stable Atmospheric Molecules
The thermodynamic properties of the oxides of sulfur
indicate that sulfur dioxide has a strong tendency to
react with oxygen in the air under normal tropospheric
conditions:
2SO2 + O2 ~ 2SO3.
(1)
Thus the ratio of [SO3]/tSO2] is about 8 x 1011 at equi-
librium in air at 1 atm and 25°C. (Square brackets
indicate the concentration of the species inside the
brackets.) Thermodynamic arguments also tell us that at
humidities normally encountered in the lower troposphere,
the SO3 produced by reaction (1) will be converted
efficiently to sulfuric acid, H2SO4(aq), according to
SO3 + H2O ~ H2SO4(aq).
(2)
The reaction of SO3 with H2O is so fast that any
process in which SO3 is formed in the moist troposphere
can be considered equivalent to the formation of H2SO4.
Certain metal ions (Mn2+, Fe3+, etc.) in aqueous
solutions of SO2(HSO5) can catalyze the overall
sequence of reactions (1) and (2). However, thermo-
dynamics tells us nothing about the rates of chemical
reactions, and the rate of reaction (1) is so slow under
catalyst-free conditions in the gas phase that it can be
neglected as a source of sulfuric acid in the atmosphere.
The thermal oxidation of NO and NO2 in the gas phase
is also slow. Pathways that are thermodynamically
favored are
OCR for page 157
157
2NO + O2 + 2NO2,
2NO2 + H2O ~ HNO3 + MONO.
(3)
(4)
Reaction (3) requires literally days for the conversion
of a significant fraction of the nitric oxide present [at
concentrations of the order of parts per billion (ppb)]
in the troposphere, and the homogeneous gas-phase reaction
(4) is immeasurably slow (Schwartz and White 1982).
Obviously, other, seemingly less direct reactions must be
invoked to account for the observed rates of SO2, NO, and
NO2 oxidation, which are between 1 and 100 percent/in.
A number of the more complex pathways involve photo-
chemistry. In one, sulfur dioxide absorbs light in the
ultraviolet region of the solar radiation incident in the
troposphere, and, in principle, excited states of SO2
generated in this fashion could lead to SO2 oxidation
in the troposphere. Figure A.1 indicates that there is a
significant overlap between the actinic flux incident in
the lower troposphere and two distinct absorption regions
of SO2. Excitation in the "forbidden," long-wavelength
band forms the excited SCt(3B1) species, while excitation
at the wavelengths below 330 nm generates higher excited
states, presumably the 1A2 and 1B1 species.
SO2(XlAl) + hv(340 < A < 400 nm) ~ So2(3Bl),
SO2(XlAl) + hv(240 < ~ < 330 nm) ~ SO2(1A2, 1Bl). (6)
These excited states of SO2 are nondissociative; only
quanta of light at wavelengths below 218 nm (which do not
penetrate to the troposphere) provide sufficient energy
to allow photodissociation:
~1
SO2(X*A1) + hv(A < 218 nm)
~ S02(1B2) ~ 0(3P) + So(3£ ) e
(7)
The lower excited singlet states of SO2(1A1, 1Bl) appear
to be very short lived in air at 1 atm, and they are
rapidly converted by collisional perturbations to S02(3B1)
molecules, possib~ SO2(3A2) and SO2(3Bk) molecules, and
ground-state S02(X A1) molecules. The rate of excitation
of SC2 through absorption of sunlight can be very sig-
nificant. If this excitation were the rate-determining
step in the photooxidation of SO2, that is, if every
molecule of SO2 that is photoexcited were oxidized
OCR for page 158
158
400
_ 300
c,
. _
cat
-
c'
to
0 200
-
c~
-
x
111
100 _
O
240
1 1 1 " ~.~ 1 -
320
Wavelenoth (nary)
0.10
0.08 A,
c
. _
. _
-
o
0.04 _
C1
0.02
400
FIGURE A.1 Comparison of the extinction coefficients (liters mole~1 cm~i, base 10)
of SO2 within the first allowed band (left), the "forbidden" band (right), and a typical
distribution of the flux of solar quanta (relative) at ground level (dashed curve).
SOURCE: Calvert et al. (l978~.
through subsequent reaction with O2 or other reactants,
then the lifetime of SO2 in the lower troposphere with
overhead sun should be as low as 52 minutes (Sidebottom
et al. 1972). Of course, this is not the case. The
SO2(3B1) species appears to be one of the most favored
states, which is ultimately populated through absorption
of sunlight and collisional processes in the lower
atmosphere. The reactions of this species with various
atmospheric gases and many atmospheric impurities have
been studied extensively (Calvert et al. 1978). Quenching
of SO2(3B1) by atmospheric gases is expected to be the
dominant process. In air at 1 atm, 25°C, and 50 percent
relative humidity, quenching by N2, 02, H2O, and Ar will
occur 45.7, 41.7, 12.2, and 0.3 percent of the time,
respectively. Quenching by impurity gases is highly
improbable. Even when impurities are present at
concentrations of the order of ppm, the rates of SO2
conversion by such species are very slow (Calvert et al.
1978). All available evidence suggests that the only
. . . . . .
significant chemical result of the major quenching
reactions of SO2(3B1) occurs with O2, and this does not
lead efficiently to any overall chemical change in the
SO2, but low-lying, excited electronic states of molecular
oxygen, O2(1Ag) and O2(1£+g), are formed by energy
transfer.
OCR for page 159
159
SO2 ( B1) + O2 ( Eg) + SO2 (X A1) + O2 ( fig)
~ SO2(2 Al) + O2( Ag)
(8)
.
(9)
These reactions are not unique sources of excited oxygen,
species that are also formed in other atmospheric
reactions at much higher rates (Calvert et al. 1978).
Thus we conclude that photooxidation of SCt is also not
an important source of acids in the atmosphere.
Similar conclusions are reached from a study of the
photochemistry of NO and NO2. NO does not absorb solar
radiation in the wavelength range available near the
Earth's surface, so its photochemistry is not important
in the troposphere. NO2 does absorb radiation over a
wide range, and absorption of the wavelengths below 430
nm leads to dissociation (NO2 + he ~ O + NO). However,
formation of acid as a direct result of the photochemistry
.
of NO and NO2 is unimportant.
Reactive Transient Species in the Troposphere
Most of the gas-phase tropospheric chemistry of SO2,
NO, NO2, and other impurity molecules involves reactions
with a variety of reactive excited molecules, atoms, and
free radicals (neutral fragments of stable molecules)
formed by absorption of sunlight by trace gases in the
atmosphere. As a background to discussions to come we
review briefly here some of this important chemistry
since it enters directly or indirectly into many of the
reaction pathways that lead to the formation of acids in
the troposphere.
In the polluted troposphere, NO2 is dissociated by
sunlight absorption (A < 430 nm) to form reactive, ground
state oxygen atoms, O(3P), and NO, while the oxygen atom
reacts rapidly to form ozone (O3):
NC: + hv(A < 430 nm) + O(3P) + NO,
O(3P) + O2(+M) ~ O3(+M).
Ozone can reoxidize NO to NC2 in (12) or react with
alkenes to give highly reactive ozonides in (13) and
Criegee intermediates in (14):
(10)
Ill)
OCR for page 160
160
0 3 + NO ~ O2 + NO2
0 - 0 - 0
O3 + RHC=CHR ~ RHC- CHR
0 - 0 - 0
1 1
RHC - CHR ~ RCHO2 + RCHO
Ozone can also oxidize NO2 to the reactive transient
NO3 in (15), and this can lead to N2O5 in (16):
O3 + NO2 + O2 + NO3,
Net + Net (+M) By N2O5 (+M) .
(12)
(13)
(14)
(15)
(16)
The photodecomposition of ozone may generate electroni-
cally excited oxygen atoms, O(1D), and excited molecular
oxygen with absorption in the short-wavelength region of
the spectrum:
O3 + hv(290-306 nm) ~ O(1D) + O2( fig),
O3 + hv(290-350 nm)
~ O(1D) + O2 or O + 02(1Ag' 3£g)'
O3 + hv(450-700 nm) ~ O + O2.
(17)
( 18)
(19)
The O(1D) species formed in (17) is much more reactive
than the ground-state oxygen atoms [O(3P), often simply
symbolized by O]. O(1D) reacts efficiently when it
collides with a water molecule to form a highly important
transient in atmospheric chemistry, the hydroxy radical,
HO:
O (ID) + H2O ~ 2HO.
(20)
This radical, unlike many molecular fragments formed from
carbon-containing molecules, is unreactive toward oxygen,
and it survives to react with most atmospheric impurities
such as the hydrocarbons, aldehydes, NO, NO2, SO2, and
CO. Its reactions with carbon monoxide and the hydro-
carbons (RH) lead to another important class of reactive
transients, the peroxy radicals:
HO + CO ~ H + CO2,
(21)
OCR for page 161
161
H + O2(+M) + HO2(+M),
HO + RH ~ H2O + R.
R + O2(+M) ~ RO2(+M).
(22)
(23)
(24)
Here R represents the alkyl groups such as methyl (CH3),
ethyl (C2H5), or another larger group derived from the
parent hydrocarbons, methane (CH4), ethane (C2H6), or
larger hydrocarbon (RH), respectively. The reaction of the
HO radical with aldehydes (RCHO) forms the acyl (RCO) and
acylperoxy (RCO02) radicals in similar reactions:
RCHO ~ HO ~ RCO + H2O,
RCO + O2(+M) ~ RC002(+M).
(25)
(26)
The peroxy radicals react rapidly with NO to form NO2 and
other classes of reactive species. In the case of the
HC2-No reaction, HO is regenerated, while with the RO2 and
SCOW radicals, alkoxy (RO) and acyloxy (RCO2) radicals,
respectively, are formed:
HO2 + NO ~ HO + NO2,
RO2 + NO ~ RO + NO2,
RCOO2 + NO ~ RCO2 + NO2
(27)
(28)
(29)
The most common fate of the smaller alkoxy radicals in the
lower atmosphere is reaction with oxygen, leading to HO2
radicals and a carbonyl compound. For example, with the
simplest alkoxy radical, methoxy (CH3O), the following
reaction occurs:
CH3O + O2 + HO2 + CH2O
(30)
The RCC: radicals are of short lifetime, decomposing to
form an alkyl radical (R) and CO2, with the subsequent
generation of another peroxyalkyl radical:
RCO2 ~ R + CO2,
R + O2(+M) ~ RO2(+M)
(31)
(32)
Reactions (14)-(26) combined with reactions (27)-(32) form
a chain reaction. That is, a single initial HO radical
OCR for page 162
162
may oxidize CO, hydrocarbon, or aldehyde, and additional
HO radicals will be regenerated as NO is oxidized to N02;
the subsequent steps occur again and again in a repeating
cycle of events. The peroxy radicals and ozone are the
principal oxidizing agents for NO in the lower atmosphere
[reactions (IS) and (27)-(29)] (Demerjian et al. 1974).
When peroxy radicals react with NO2, an additional
class of highly reactive compounds is generated--the
peroxynitrates:
HO2 + NC2 ~ HC2NC2,
RO2 ~ NO2 ~ RO2NO2 ~
RCOO2 + NO2 ~ RCOO2 NO2
Peroxynitric acid formed in (33), and the alkyperoxy-
nitrates formed in (34) are relatively unstable in the
lower troposphere at temperatures common in summer
months; they dissociate readily to reform peroxy radicals
and NC2. However, during the cold winter months or in
the stratosphere they can act as temporary sinks for
HO2 and RO2 radicals and NO2. Peroxyacylnitrates
(RC ~ NO2), of which peroxyacetylnitrate (CH3COCtNC2) is
the most common, have longer lifetimes and can be the
source of radical generation even during the nighttime
hours.
The excited singlet delta molecular oxygen, O2(1Ag),
product of reactions (17) and (18), and excited singlet
sigma oxygen, O2(1~+)' the product of reaction (8), can
also be created by ~irect absorption of sunlight by
atmospheric O2 and by energy transfer reactions from
other photoexcited species such as NO2(\ > 430 nm) and
excited triplet aromatic hydrocarbons.
(33)
(34)
(35)
For current purposes the complex array of interactions
that occur among the reactive species outlined here and
with the various atmospheric impurities need not be
considered. It suffices to say that many aspects of
tropospheric chemistry, including photochemical "smog"
and ozone generation, depend on these happenings
(Demerjian et al. 1974). It is important, however, to
evaluate the potential significance of the many highly
reactive transient species of the atmosphere for reactions
that oxidize SO2 and NO2 to acids.
OCR for page 163
163
Atmospheric Oxidation of SO2
by Reactive Transient Species
There are a large number of potentially significant
gas-phase reactions of the reactive transients leading to
oxidation of SC: in the troposphere. The potential
candidate reactions, summarized in Table A.1, have the
thermodynamic potential to occur as measured qualitatively
by the sign of the change in enthalpy (aH°) for the
overall reaction. Rate constants for most of these
elementary reactions have been determined. These data,
coupled with estimates of the concentrations of the
transients in the atmosphere, allow us to evaluate the
significance of each reactant in oxidizing SO2 in the
atmosphere. Such evaluations have shown that reactions
(42), (50), (56), and (59) or (61) are potentially sig-
nificant sources of SO2 oxidation; some of these
reactions are only important for certain peculiar
atmospheric conditions. By far the most important of the
gasphase reactions is the reaction of HO radicals with
so2
HO + SO2(+M) ~ HOSO2(+M)
.
(56)
The rate-constant data for reaction (56) have been
reviewed recently (Calvert and Stockwell 1983) and are
summarized in Figure A.2. The theoretical effect on this
rate constant of the altered temperature and pressure of
the atmosphere at various altitudes is shown in Figure
A.3.
The "best" values of kII(56) suggested from the
analysis given by Calvert and Stockwell (1983) are
somewhat larger than those chosen by Moortgat and Junge
(1977), Zellner (1978), and the values recommended for
use by the CODATA (1980) group and the NASA (1979) panel
and are more consistent with currently available infor-
mation on the HO-SO2 reaction. One must, however,
retain considerable pessimism about the accuracy of all
these data; realistic confidence limits should include
+50 percent of the suggested value.
Evidence is good that HOSO2 formed in reaction (56)
ultimately leads to the generation of sulfuric acid
aerosol. However, HOSO2 is not a stable molecule; it
is a free radical that is probably highly reactive with
several atmospheric compounds. It is not now clear what
elementary reaction
conversion
to H2SO4.
pathways are important in its
Although there has been a great
OCR for page 164
164
TABLE A. I Enthalpy Changes and Recommended Rate Constants for Potentially
Important Reactions of Ground State SO~ and SO~ Molecules in the Lower
Atmosphere
-
-~H ~kh
Reaction (kcal/mole) (25°C) (cm3/molec-sec)
(36) O. ('~g) + SO. ~ SO4 (biradical; cyclic) ~25, ~28
(37) O. ('Ax) + SO. ~ SO,3-O (3P) -13.5 3.9 x l0-2°
(3X) O. (~Ag) + SO.-O. (~5g) + SO. 22.5
(39) O. (~5g) + SO2-SO4 (biradical: cyclic) ~40, ~43
(40) O. ('Ig) + SO.-SO~ + O (3P) 1.5 6.6 x 10-~6
(41) O2 (~g) + SO.-SO. + O. ('/`,~) 15.0
(42) O (3P) + SO~ (+M)-SO~ (+M) 83.0 5.7 x 10-~4
(43) O3 + SO2-O. + SO; 57.6 <8 x 10-24
(44) NO2 + SO2-NO + SO3 9.9 8.8 x 10-3°
(45) NO,3 + SO2-NO2 + SO3 32.6 <7 x 10-2'
(46) ONOO + SO.-NO. + SO,3 ~30 <7 x 10-2'
(47) N2O~ + SO2 ~ N.O4 + SO3 24.0 <4 x 10-23
(48) HO. + SO'-HO + SO3 16.7 )
<1 x 10-'8
(49) HO. + SO. (+M)-HO.SO2 (+M) ~7
(50) CH3O2 + SO2-CH3O + SO3 ~27 <1 x 10-'8
(51) CH,;02 + SO2 (+M) _ CH3O2SO2 (+M) ~31 ~1.4 x 10-'4C
(52) (CH,3),~CO2 + SO2-(CH3)3CO + SO,3 ~26 /
<7.3 x 10-~9
(53) (CH~),~CO2 + SO2 _ (CH,~)3CO2SO2 ~30
(54) CH,~COO. + SO2-CH3CO. + SO,1 ~33
<7 x 10-19
(55) CH,~COO2 + SO2-CH3COO,SO2 ~37
(56) HO + SO2 (+M) ~ HOSO2 (+M) ~37 1.1 x lo-~2
(57) CH30 + SO2 (+M)_ CH3OSO2 (+M) ~24 ~5.5 x 10-13
Q~ ,0
(58) RCH-CHR + SO2 ~ 2RCHO + SO3 ~69 See text
Q o,-O
RCH-CHR + SO2-2RCHO + SO3 ~89 See text
(59) RCHOO- + SO2_ RCHO + SO,~ ~79 k59/k60~ 6 x 10-5
(60) RCHOO- + H2O-RCOOH + H2O ~121 (R = CH3)
(61) RCHO ~ SO2-RCHO + SO3 ~58 k6l/k62 ~ 4
Q Q
(62) RCHO + CH2O ~ RCHOCH2O ~12
(63) SO3 + H2O ~ H2SO4
24.8
(R = H)
, ,, 9.1 x 10
'Enthalpy change estimates were derived from the data of Benson (1978), Harding and Goddard (1978),
and Domalski ( 1971).
~Rate constants are expressed as second-order reactions for I atm of air at 25°C; see Calvert and
Stockwell (198,3) for the references to the original literature.
''The reverse reaction is so fast that the rate of oxidation of SO2 via (51) is very dependent on alternate
fates of the CH3O2SO2 species.
OCR for page 165
165
1 000
E
~ 00
-
10
~ 1 ~ 1
1
0.01 0.1
_ I 1--l rilill r -Al l-:lil --I 1-l 1 IrTl ~-r I I Al ITT:
C Corrected data of Castleman and Tang (1976-1977)
O Leu (1982)
a Range of values reported by Harris and Wayne (1975)
Cox and Sheppard (1980)
O Davis et a I. ( 1979)
D' ,/
o
i I I I l_L
/
1/[M] (cm3/molec x 1018)
3'
new ,
MU
_4
~ J /^
/
/
/
I 1 1 1 1 1 11 1 1 1 1 1 1 1 1 1 1 1
1.0 1n 100
10
FIGURE A.2 Comparison of the experimental data for the effective second-order
rate constants for the reaction (56) with M = N2. SOURCE: Calvert and Stockwell
(1983~.
amount of speculation in this regard, there is little
experimental evidence to help determine the relative
importance of suggested alternative routes. Thus Cox
(1974-1975, 1975), Calvert and McQuigg (1975), Calvert et
al. (1978), Davis et al. (1979), Friend et al. (1980),
Leu (1982), Benson (1978), and others have suggested that
the HOSO2 radical may participate in a variety of
radical-radical and radical-molecule reactions. These
reactions are summarized in Figure A.4. Although the
mechanism of H2SO4 generation following (56) is unclear
today, it is probable that reaction (56) is the rate-
determining step in the sequence. Recent evidence
suggests that the concentration of HO in photooxidizing
mixtures of MONO, NO, NO2, and CO is insensitive to
even large additions of SC2 (Stockwell and Calvert
1983). Thus the following sequence of reactions seems
favored:
OCR for page 191
191
104
103
1o2
-
~ 10
a)
Q
-
LL 1
CC
A
O 10
0 10 2
10-3
10-4
10-5
H 2O2
M//
/ /
/ Mn:
/NO2
0 1 2
/NO2
3 4 5 6
Gas-Phase Concentrations (ppb)
H2O2 1
O3 50
HNO2 2
NO2 1
Liquid-Phase Concentrations
Fe3+ 3 x 10-7 M
Mn2+ 3 x 10~ M
C 1x10~2g/liter
pH
FIGURE A.13 Theoretical rates of liquid-phase oxidation of SO2 assuming 5 ppb of
SO2, 1 ml/m3 of water in air, and concentrations of impurities as shown. SOURCE:
Martin (1983~.
For the hydrated NO2 oxidation, Martin used the
Henry's law constant and rate constants of Schwartz and
White (1982) and Lee and Schwartz (1981). The concen-
trations of NO2, MONO, and NO are related thermo-
dynamically, and so they may not be regarded as
completely independent.
The trend of rising conversion rates with increased pH
results from either the rising equilibrium concentration
of sulfur (IV) or the sensitivity of the rate constants
to pH, or both. The relative conversion rate increases
with rising concentrations of sulfur (IV) in this range
of pH values and for this amount of liquid. H2O2, which
OCR for page 192
192
undergoes an acid-catalyzed reaction, is the only oxidant
for which the rate dependence on [H+] compensates for
the decreased solubility of SO2 with increased [H+];
hence the rate of oxidation by H2O2 is nearly
independent of pH.
Martin emphasized that the relative positions of the
curves in Figure A.13 differ for differing assumptions
about the composition of the droplets. For example, an
aerosol with less liquid water will, as a rule, have
higher concentrations of nonvolatile species such as
carbon, iron, and manganese but a smaller reaction
volume, so the relative positions of the curves may
differ in this situation. The relative positions will
also be different at night, when the concentration of
photochemically derived oxidants drops. At temperatures
below 25°C, the rate constants are lower in accordance
with the activation energies of the given reactions, and
the Henry's law coefficients are higher. The two effects
act in opposite directions. In some cases, such as that
for H2O2 oxidation, the net rate rises as the temperature
falls (at constant gas-phase concentration). In other
cases, such as for iron catalysis, the net rate falls
with temperature.
The concentrations of reactants used in deriving
Figure A.13 are representative of those that might be
anticipated in the atmosphere. Oxidation by H2O2
dominates all reactions for conditions of low pH.
Oxidation rates can be greater than 100 percent/in. The
rate for oxidation by ozone varies from about 10
percent/in at pH 4.5 to about 1 percent/in at pH 4.0. The
contributions from the Fe3+, Mn/+, and carbon-catalyzed
reactions (for the conditions specified) are below 1
percent/in for solutions of pH ~ 4.5. Oxidation by NO2
and HONO is even less significant under these conditions.
An additional influence on the rates and kinetics of
sulfite oxidation in clouds and rainwater can arise from
the complexation of HSOi(aq) by aldehydes scavenged by
the droplets. The common gaseous products of atmospheric
oxidation of the impurity hydrocarbons, CH2O, CH3CHO,
CH3COCH3, and possibly other carbonyl compounds, can in
principle complex with HSO3. The result of such
interactions could increase the solubility of SO2 in
the droplet and conceivably retard sulfite oxidation by
the oxidants.
Although there is significant theoretical evidence
that H2O2 may be the most important oxidizing agent for
acid generation in cloud water and rain, unambiguous
OCR for page 193
193
experimental measurements of H2O2 levels in air and in
cloud water and rain have not been possible to date.
Peroxide development in the sampling train and lack of
selectivity of the luminol technique employed in previous
work prevents a firm conclusion about the H2O2 levels in
the atmosphere and in precipitation (Heikes et al. 1982,
Zika and Saltzman 1982). Theoretical simulations of the
atmospheric chemistry of a mixture of reactants in a
highly diluted urban atmosphere show that H2Ok generation
through reactions (71) and (73) can provide substantial
levels of H2O2, CH302H, and other hydroperoxides in the
gas phase. If these species are absorbed i non In
aerosols, cloud water, or rainwater, for example, as is
probable for H2O2 in view of its very high Henry's law
constant, then oxidation by H2O2, and possibly other
peroxides, can be most significant. The possible roles
for peroxyacetylnitrate, peroxynitr~c acid, CH3C2H,
and other peroxides in solution-phase sulfite oxidation
remain to be evaluated.
Solution-Phase Generation of Nitric and
Nitrous Acids in the Troposphere
There is some evidence of the formation of HONO2 in
clouds and rainwater. Recently, both theory and experi-
ment suggest that HONO2 may be formed rapidly from a
combined gasphase/liquid-phase process. Through
simulation, Heikes and Thompson (1981) have suggested
that N205 generated by O3-NO2 reactions (15) and (16) may
be scavenged effectively by H2O droplets to form HONO2
in clouds or rainwater.
NO2 + O3 ~ O2 + NO3,
N ~ ~ NO>(+M) ~ N2O5(+M),
N2O5 + H2O(liq) ~ 2H+(aq) ~ 2NO3(aq)
(15)
(16)
(81)
A preliminary report of the experimental observation of
this process has been made recently by Gertler et al.
(1982). An evaluation of the contribution of this
mechanism to the total [HONO2] found in the atmosphere is
not now possible.
Lee and Schwartz (1981) studied the rates of reactions
(82) and (83):
OCR for page 194
194
2NO2(g) + HzO(liq) ~ 2H+(aq) + NO3(aq) + NO2(aq), (82)
NO(g) + NO2(g) + H2O(liq) ~ 2H+(aq) + 2NO2(aq). (83)
They found that at the lower partial pressures character-
istic of moderately polluted atmospheres the rates are
slow (10-9 to 10-8 M/h). Unless high partial pressures
(e.g., 10-7 atm) of these gases are maintained in contact
with liquid water for substantial periods of time (tens of
hours), reactions (82) and (83) cannot represent a substan-
tial source of atmospheric acidity. Lee and Schwartz found
no evidence of Fez+ ion catalysis of these reactions, but
they suggested that possibly other metal ions could enhance
these rates. There is no evidence related to this pos-
sibility now available.
Although significant uncertainty remains concerning the
source of HNO3 in clouds and rainwater, the limited
evidence currently available favors the probable importance
of the formation of N2O5 in (15) and (16) followed by its
reaction in water droplets to form HNO3.
SUMMARY
In summary, a wide variety of interrelated homogeneous
gas-phase, solution-phase, and heterogeneous chemistry may
result ultimately in oxidation of SO2 to sulfuric acid and
NOX to nitric acid. The homogeneous gas-phase oxidation of
SC2 by the HO radical and the solution-phase oxidation of
S(IV) through H2O2, O3, and possibly other species appear to
be the major sources of H2SO4. In the cloud-free,
ambient, sunlight-irradiated troposphere, nitric acid is
probably generated largely by the reaction of HO radicals
with NO2. Both HONO2 and H2SO4 produced in the gas phase
can be scavenged effectively by cloud water and precipita-
tion. NO2 may be oxidized to HONO2 if sufficient O3 and
NO2 are present. Following its gas phase generation, N2O5
may be scavenged effectively by water droplets to form
HONO2.
All the various pathways that lead to the oxidation of
SO2 and NOX are coupled by common products and reactants
that can directly and indirectly influence the rates of
reaction by the other pathways. In general the homogene-
ous gas-phase reactions can lead to maximum daylight
rates of acid formation of a few percent per hour for
SO2 and 20-30 percent/in for NOk. Solution-phase
reactions involving H2O2 and O3 can in principle
OCR for page 195
195
convert SO2 to H2SO4 in cloud water and precipitation at
much higher rates (as high as 100 percent/in) for concen-
trations Of H2O2 and O3 in the troposphere that
reasonably could result from the normal homogeneous
reactions characteristic of atmosphere chemistry during
daylight.
REFERENCES
Anderson, L.G. 1980. Absolute rate constants for the
reaction of OH with NO2 in N2 and He from 225 to
389K. J. Phys. Chem. 84:2152-2155.
Aubuchon, C. 1976. The rate of iron catalyzed oxidation
of sulfur dioxide by oxygen in water. Ph.D. thesis.
The Johns Hopkins University, Baltimore, Md.
Barrie, L.A., and H.W. Georgii. 1976. An experimental
investigation of the absorption of sulfur dioxide by
water drops containing heavy metal ions. Atmos.
Environ. 10:743-749.
Benson, S.W. 1978. Thermochemistry and kinetics of
sulfur-containing molecules and radicals. Chem. Rev.
78:23-35.
Brimblecombe, P., and D.J. Spedding. 1974. The reaction
order of the metal ion catalyzed oxidation of sulfur
dioxide in aqueous solution. Chemosphere 1:29-32.
Brodzinski, R., S.G. Chang, S.S. Markowitz, and T.
Novakov. 1980. Kinetics and mechanism for the
catalytic oxidation of sulfur dioxide on carbon in
aqueous suspension. J. Phys. Chem. 84:3354-3358.
Calvert, J.G., and R.D. McQuigg. 1975. The computer
simulation of rates and mechanisms of photochemical
smog formation. Int. J. Chem. Kinet. Symp. 1:113-154.
Calvert, J.G., and W.R. Stockwell. 1983. The mechanism
and rates of the gas phase oxidations of sulfur
dioxide and nitrogen oxides in the atmosphere. In Acid
Precipitation: SO2, NO, and NO2 Oxidation
Mechanisms: Atmospheric Considerations. Ann Arbor,
Mich.: Ann Arbor Scientific Publications. In press.
Calvert, J.G., F. Su, J.W. Bottenheim, and O.P. Strausz.
1978. Mechanism of the homogeneous oxidation of sulfur
dioxide in the troposphere. Atmos. Environ. 12:197-226.
Campbell, M.J., J.C. Sheppard, and B.F. Au. 1979.
Measurements of hydroxyl concentration in the boundary
layer air by monitoring CO oxidation. Geophys. Res.
Lett. 6:175-178.
OCR for page 196
196
Cass' G.R., and F.H. Shair. 1983. Sulfate accumulation in
a sea breeze/land breeze circulation system. J.
Geophys. Res. In press.
Castleman, A.W., Jr., and I.N. Tang. 1976-1977. Kinetics
of the association reaction of SO2 with hydroxyl
radical. J. Photochem. 6:349-354.
Chameides, W.L., and D.D. Davis. 1982. The free radical
chemistry of cloud droplets and its impact upon the
composition of rain. J. Geophys. Res. 87:4863-4877.
Chang, J.S., P.J. Wuebbles, and D.D. Davis. 1977. A
theoretical model of global tropospheric OH distribu-
tions. UCRL-78392. Livermore, Calif.: Lawrence
Livermore Laboratory. 22 pp.
Chang, S.G., R. Brodzinsky, R. Toossi, S.S. Markowitz,
and T. Novakov. 1978. Catalytic oxidation of SC2 on
carbon in aqueous suspensions. Pp. 122-130 in
Conference on Carbonaceous Particles in the
Atmosphere. Berkeley, Calif.: Lawrence Berkeley
Laboratory.
Chang, S.G., R. Roost, and T. Novakov. 1981. The
importance of soot particles and nitrous acid in
oxidizing SO2 in atmospheric aqueous droplets.
Atmos. Environ. 15:1287-1292.
CODATA. 1980. Evaluated kinetic and photochemical data in
atmospheric chemistry. J. Phys. Chem. Ref. Data
9:295-471.
Coughanowr, D.R., and R.E. Krause. 1965. The reaction of
SO2 and O2 in aqueous solutions of MnSO4. Ind.
Eng. Chem. Fundam. 4:61-66.
Cox, R.A. 1974-1975. The photolysis of nitrous acid in
the presence of carbon monoxide and sulfur dioxide. J.
Photochem. 3:291-304.
Cox, R.A. 1975. The photolysis of gaseous nitrous acid--a
technique for obtaining kinetic data on atmospheric
photooxidation reactions. Int. J. Chem. Kinet. Symp.
1:379-398.
Cox, R.A., and D. Sheppard. 1980. Reaction of OH radicals
with gaseous sulfur compounds. Nature 284:330-331.
Crutzen, P.J., and J. Fishman. 1977. Average concentra-
tions of OH in the troposphere and the budgets of
methane, carbon monoxide, molecular hydrogen and
1, 1,1-trichloromethane. Geophys. Res. Le~t. 4:321-324.
Davis, D. D., A. R. Ravishankara, and S. Fischer. 1979.
SO2 oxidation via the hydroxyl radical: atmospheric
fate of HSOX radicals. Geophys. Res. Lett. 6:113-116.
Demerjian, K.L., J.A. Kerr, and J.G. Calvert. 1974. The
mechanism of photochemical smog formation. Adv.
Environ. Sci. Technol. 4:1-262.
OCR for page 197
197
Domalski, E.S. 1971. Thermochemical properties o f
peroxyacetyl (PAN) and peroxybenzoyl nitrate (PBN).
Environ. Sci. Technol. 5:443-444.
Ea tough, D.J., B.E. Richter, N.L. Ea tough, and L.D.
Hansen. 1981. Sulfur chemistry on smelter and powe r
plant plumes in the western U.S. Atmos. Environ.
15:2241-2253.
Er ickson, R.E., L.M. Yates, R.L. Clark, and D. McEwen.
1977. The reaction of sulfur dioxide with ozone in
water and its possible atmospheric significance.
Atmos. Environ. 11:813-817.
Forrest, J., R.W. Garber, and L. Newman. 1981. Conversion
rates in power plant plumes based on filter pack
data: the coal-fired Cumberland plume. Atmos.
Environ. 15:2273-2282.
Freiberg, J.E., and S.E. Schwartz. 1981. Oxidation of
SO2 in aqueous droplets: mass-transport limitation
in laboratory studies and the ambient A~mC,c:nh~r"
Atmos. Environ. 15:1145-1154.
~.. _, ~ -
Fr lend, J.P., R.A. Barnes, and R.M. Vasta. 1980.
Nucleation by free radicals from the photooxidation of
sulfur dioxide in air. J. Phys. Chem. 84:2423-2436.
Fuzzi, S. 1978. Study of iron (III) catalyzed sulfur
d ioxide oxidation in aqueous solution over a wide
range of pH. Atmos. Environ. 12:1439-1442.
Garber, R.W., J. Forrest, and L. Newman. 1981.
Conversion
rates in power plant plumes based on filter pack
data: the oil fired Northpor~ plume. Atmos. Environ.
15:2283-2292.
Gertler, A.W., D.F. Miller, D. Lamb, and U. Katz. 1982.
SO2 and NO2 reactions in cloud droplets. Pp.
112-115 in Abstracts, American Chemical Society, 185th
National Meeting, Las Vegas, Nev.
Gillani, N.V., and S. Kohli. 1981. Gas-to-particle
conversion of sulfur in power plant plumes. I.
Parameterization of the conversion rate for dry,
moderately polluted ambient conditions. Atmos.
Environ. 15:2293-2313.
Graedel, T.E., L.A. Farrow, and T.A. Wicker. 1976.
Kinetic studies of the photochemistry of the urba n
troposphere. Atmos. Environ. 10:1095-1117.
Harding, L.B., and W.A. Goddard III. 1978. Mechanism of
gas phase and liquid phase ozonolysis. J. Am. Chem.
Soc. 100:7180-7188.
Harris, G.W., and R.P. Wayne. 1975. Reaction of hydroxyl
radicals with NO, NO2, and SO2. J. Chem. Soc.,
Faraday Trans. 1. 71:610-617.
OCR for page 198
198
Hecht, T.A., and J.H. Seinfeld. 1972. Development and
validation of generalized mechanism for photochemical
smog. Environ. Sci. Technol. 6:47-57.
Hegg, D.A., and P.V. Hobbs. 1980. Measurements of
gas-to-particle conversion in the plumes from five
coal-fired electric power plants. Atmos. Environ.
14:99-116.
Heikes, B.G., and A.M. Thompson. 1981. Nitric acid
formation in-cloud gas phase, photochemical and/or
heterogeneous paths. Eos 62:884.
Heikes, B.G., A.L. Lazrus, G.L. Kok, S.M. Kunen, B.W.
Gandrud, S.N. Gitlin, and P.D. Sperry. 1982. Evidence
for aqueous phase hydrogen peroxide synthesis in the
troposphere. J. Geophys. Res. 87:3045-3051.
Henderson, R., and K. Wingartner. 1980. Analysis of MAP3S
Precipitation Chemistry Data. MTR-8~W00265. McLean,
Va.: MITRE Corporation.
Hoather, R.C., and C.F. Goodeve. 1934a. The oxidation of
sulphurous acid. I. The dilatometric technique. Trans.
Faraday Soc. 30:626-629.
Hoather, R.C., and C.F. Goodeve. 1934b. The oxidation of
sulphurous acid. III. Catalysis by manganese sulfate.
Trans. Faraday Soc. 30:1149-1156.
Hoffman, M.R., and J.O. Edwards. 1975. Kinetics of the
oxidation of sulfite by hydrogen peroxide in acidic
solution. J. Phys. Chem. 79:2096-2098.
Hoffman, M.R., and D.J. Jacob. 1983. Kinetics and
mechanisms of catalytic oxidation of dissolved sulfur
dioxide in aqueous solution: an application to
nighttime fog water chemistry. In Acid Precipitation:
SO2, NO, NO2 Oxidation Mechanisms: Atmospheric
Considerations. Ann Arbor, Mich.: Ann Arbor Science
Publishers, Inc. In press.
Hoffman, M.R., S.D. Boyce, A. Hong, and L. Moberly. 1982.
Catalysis of the autooxidation of equated sulphur
dioxide by homogeneous and heterogeneous transition
metal complex. In Heterogeneous Catalysis: Its
Importance in Atmospheric Chemistry, D. Schryer, ed.
AGU Monograph. Washington, D.C.:
Union.
American Geophysical
Holt, B.D., P.T. Cunningham, and R. Kumar. 1981. Oxygen
isotopy of atmospheric sulfates. Environ. Sci.
Technol. 15:804-808.
Holt, B.D., R. Kumar, and P.T. Cunningham. 1982. Primary
sulfates in atmospheric sulfates: estimation of
oxygen isotope ratio measurement. Science 217:51-53.
OCR for page 199
199
r
Holt, B.D., P.T. Cunningham, A.G. Engelkemeir, D.G.
Graczyk, and R. Kumar. 1983. Oxygen-18 study of
nonaqueous-phase oxidation of sulfur dioxide. Atmos.
Environ. 17:625-632.
Hov, O., and T.S.A. Isaksen. 1979. Hydroxy and peroxy
radicals in the polluted tropospheric air. Geophys.
Res. Lett. 6:219-222.
Kan, C.S., J.G. Calvert, and J.H. Shawl 1981. Oxidation
of sulfur dioxide by methylperoxy radicals. J. Phvs.
Chem. 85:1126-1132.
Larson, T.V., N.R. Horike, and H. Harrison. 1978.
~.. ~.
Oxidation of sulfur dioxide by oxygen and ozone in
aqueous solution: a kinetic study with significance
to atmospheric rate processes. Atmos. Environ.
12:1597-1611.
Lazrus, A., P. L. Haagenson, G. L. Kok, C. W. Kreitzberg,
G.E. Likens, V.A. Mohnen, W.E. Wilson, and J.W.
Winchester. 1983. Acidity in air and water in frontal
precipitation. Atmos. Environ. 17:581-591.
Lee, Y.N., and S.E. Schwartz. 1981. Reaction kinetics of
nitrogen dioxide with water at low partial pressure.
J. Phys. Chem. 85:840-848.
Leu, M.-T. 1982. Rate constants for the reaction of OH
with S°k at low pressures. J. Phys. Chem.
86:4558-4562.
Levy, H. 1974. Photochemistry of the troposphere. Acta
Photochem. 9:369-375.
Maahs, H.G. 1982. The importance of ozone in the
oxidation of sulfur dioxide in nonurban tropospheric
clouds. In Second Symposium on the Composition of the
Nonurban Troposphere, Williamsburg, Va. Boston,
Mass.: American Meteorological Society.
Mader, P.M. 1958. Kinetics of the hydrogen peroxide-
sulfite reaction in alkaline solution. J. Am. Chem.
Soc. 80:2634-2639.
Martin, L.R. 1983. Kinetic studies of sulfite oxidation
in aqueous solutions. In Acid Precipitation: SO2,
NO, and NO2 Oxidation Mechanisms: Atmospheric
Considerations. Ann Arbor, Mich.: Ann Arbor
Scientific Publications. In press.
Martin, L.R., and D.E. Damschen. 1981. Aqueous oxidation
of sulfur dioxide by hydrogen peroxide at low pH.
Atmos. Environ. 15:1615-1621.
Martin, L.R., D.E. Damschen, and H.S. Judeikes. 1981. The
reactions of nitrogen oxides with SC2 in aqueous
aerosols. Atmos. Environ. 15:191-195.
OCR for page 200
200
Matteson, M.J., W. Stober, and H. Luther. 1969. Kinetics
of the oxidation of sulfur dioxide by aerosols of
manganese sulfate. Ind. Eng. Chem. Fundam. 8:667-687.
McMurry, P.H., and D.J. Rader. 1981. Studies of aerosol
formation in power plant plumes. I. Growth laws for
secondary aerosols in power plant plumes:
implications for chemical conversion mechanisms.
Atmos. Environ. 15:2315-2327.
Meagher, J.F., L. Stockburger III, R.J. Bonanno, and M.
Luria. 1981. Cross-sectional studies of plumes from
partially SO2-scrubbed power plant. Atmos. Environ.
15:2263-2272.
Moortgat, G.K., and C.E. Junge. 1977. The role of SO2
oxidation for the background stratospheric layer in
the light of new reaction rate data. Pure Appl.
Geophys. 115:759-774.
Morris, E.D., and H. Niki. 1973. Reaction of dinitrogen
pentoxide with water. J. Phys. Chem. 77:1929-1932.
NASA. 1979. The Stratosphere--Present and Future. R.D.
Hudson and E.I. Reed, eds. Report 1049. Greenbelt,
Md.: National Aeronautics and Space Administration.
Newman, L. 1981. Atmospheric oxidation of sulfur
dioxide: a review as viewed from powerplant and
smelter plume studies. Atmos. Environ. 15:2231-2239.
Neytzell-de Wilde, F.G., and L. Taverner. 1958.
Experiments relating to the possible production of an
oxidizing acid leach liquor by autooxidation for the
extraction of uranium. Pp. 303-317 in Proceedings of
the Second U.N. International Conference on the
Peaceful Uses of Atomic Energy, vol. 3. Geneva:
United Nations.
Niemann, B.L. 1982. Analysis of wind and precipitation
data for assessment of tranaboundary transport and
acid deposition between Canada and the United States.
Pp. 35-36 in Abstracts, American Chemical Society,
185th National Meeting, Las Vegas, Nev.
Niki, H., E.E. Daby, and B. Weinstock. 1972. Mechanisms
of smog reactions. Chapter 2 in Photochemical Smog and
Ozone Reactions, Adv. Chem. Ser. 113:16-57.
Oblath, S.B., S.S. Markowitz, T. Novakov, and S.G. Chang.
1981. Kinetics of the formation of hydroxylamine
disulfonate by reaction of nitrite with sulfites. J.
Phys. Chem. 85:1017-1021.
Penkett, S.A., B.M.R. Jones, K.A. Brice, and A.E.J.
Eggleton. 1979. The importance of atmospheric ozone
and hydrogen peroxide in oxidizing sulfur dioxide in
cloud and rainwater. Atmos. Environ. 13:123-137.
OCR for page 201
201
Perner, D., D.H. Ehhalt, H.W. Patz, V. Platt, E.P. Roth,
and A. Volz. 1976. OH radicals in the lower
troposphere. Geophys. Res. Lett. 3:466-468.
Schwartz, S.E., and J.E. Freiberg. 1981. Mass-transport
limitation to the rate of reaction of gases in liquid
droplets: Application to oxidation of SC: in
aqueous solutions. Atmos. Environ. 15:1129-1144.
Schwartz, S.E., and W.H. White. 1982. Kinetics of
reactive dissolution of nitrogen oxides into aqueous
solution. Adv. Environ. Sci. Technol. 12.
Shaw, R.W., and R.J. Paur. 1983. Measurements of sulfur
in gases and particles during sixteen months in the
Ohio River Valley. Atmos. Environ. In press.
Sidebottom, H.W., C.C. Badcock, G.E. Jackson, J.G.
Calvert, G.W. Reinhardt, and E.K. Damon. 1972.
Photooxidation of sulfur dioxide. Environ. Sci.
Technol. 6:72-79.
Spicer, C.W. 1982. Nitrogen oxide reactions in the urban
plume of Boston. Science 215:1095-1097.
Stockwell, W.R., and J.G. Calvert. 1983. The mechanism of
the HO-SO2 reaction. Atmos. Environ. In press.
Valley, S.L. 1965. U.S. Standard Atmosphere. In Handbook
of Geophysics and Space Environment. New York:
McGraw-Hill.
Wang, C.C., L.I. Davis, Jr., P.M. Selzer, and R. Munos.
1981. Improved airborne measurements of OH in the
atmosphere using the technique of laser induced
fluorescence. J. Geophys. Res. 86:1181-1196.
Williams, D.J., J.N. Carras, J.W. Milne, and A.C. Heggie.
1981. The oxidation and long-range transport of sulfur
dioxide in a remote region. Atmos. Environ.
15:2255-2262.
Wilson, J.C., and P.H. McMurry. 1981. Studies of aerosol
formation in power plant plumes. I. Secondary aerosol
formation in the Navajo generating station plume.
Atmos. Environ. 15:2329-2339.
Zak, B.D. 1981. Lagrangian measurements of sulfur dioxide
to sulfate conversion rates. Atmos. Environ.
15:2583-2591.
Zellner, R. 1978. Recombination reactions in atmospheric
chemistry. Ber. Bunsenges Phys. Chem. 82:1172-1179.
Zika, R.G., and E.S. Saltzman. 1982. Interaction of ozone
and hydrogen peroxide in water: implications for
analysis of H202 in air. Geophys. Res. Lett.
9:231-234.
Representative terms from entire chapter:
gas phase
