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Heterogeneous SO2-oxidation in the droplet phase
HETEROGENEOUS SO,-OXIDATION DROPLET PHASE
IN THE
S. BEILKE Umweltbundesamt:
Pilotstation
Frankfurt. FeldbergstraDe 45, 6000 Frankfurt a.M.. W. Germany
and G. GRAVENHORX Institut fiir Atmosphiirische Chemie, Kemforschungsanlage Jiilich, W. Germany Abstract-In order to determine the rate controiling step for sulfate formation in a heterogeneous droplet system. Sot-transfer within the gasphase towards and within droplets are calculated. equilibrium between SO2 in the gasphase and sulfur (IV) in cloud and fog droplets is reached within less than 1 second. Oxidation of sulfur (IV) to sulfate in the droplet phase proceeds slower by orders of magnitutde. Three mechanisms of SO,-oxidation are discussed: {a) SO,-oxidation by Ot in the absence of catalysts; (b) SO,-oxidation by O2 in the presence of catalysts, and (cc) SO,-oxidation by strongly oxidizing agents. Mechanism (a) contributes only to a negligible extent to sulfate formation in droplets even in the presence of typical concentrations of ammonia. Indications are strong that the major function of the SO,-NH,-&O-system is not the oxidation of SO2 to sulfate. Oxidation mechanism (b) may contribute to a significant extent to sulfate formation in urban fogs in which case the concentrations of catalysts can be sufficiently high. For clouds in remote areas with much lower catalyst concentrations SO,-oxidation by mechanism (b) seems to be of little importance. The oxidation by strongly oxidizing agents (mechanism (c)) appears to be the dominant mechanism although some experimental discrepancies have to be resolved.
1. lNTRODtJCllON There is strong evidence that the precinct part of atmospheric sulfate found in the Federai Republic of Germany is generated by oxidation of atmospheric SO, and not directly emitted as sulfate containing particles. The oxidation proceeds by 2 processes: (I) Homogeneous SO,-gasphase oxidation. (2) Heterogeneous SOI-oxidation in atmospheric droplets and on aerosol particles This paper deals with the SOIoxidation in the heterogeneous SO,-air-droplet-systern. It is attempted to integrate some of the reported experimental and theoretical investigations into a unified picture and to assess some quantitative data on sulfate formation and SO,-removal in the atmosphere. In a real atmosphere we are concerned with a heterogeneous SOI-oxidation in the aqueous phase whereas in laboratory experiments both homogeneous and heterogeneous aqueous systems were used to study SO,-oxidation. In order 10 apply the laboratory results obtained in a homogeneous aqueous phase to the atmospheric heterogeneous system some estimates are made of SO,-transport rates within both the gaseous and the droplet phase. These calculations are described in the first part of this paper. In the second part. three mechanisms of SOz-oxidation will be discussed and their importance for atmospheric sulfate formation assessed: (a) SO,-oxidation * Sulfur (IV) stands for the sum of SO,-H20, and SO:-.
HSO;
by O2 in the absence of catalysts, (b) SO1-oxidation by O2 in the presence of catalysts and (c) Sot-oxidation by strongly oxydizing agents. On the basis of this discussion further studies to 611 knowledge gaps will be suggested. 2 SO,-TRANSFER AND ARSORPTION
In a heterogeneous system in which gaseous SO2 is in equilibrium with sulfur (IV)* in the aqueous phase, the following equations describe the equilib rium, (SO,),,, + Hz0 &
SOz+I,O
SO~+z*o AH+ 7 HSO; KI,H+
(1)
-f HSO;
(2)
+ SO:-
(3)
where SOz*H20 is physically-halved SOz, HSO; is the bisulfite ion, and SOi- is the sulfite ion. H = Henry constant; K, = first dissociation constant; K, = second dissociation constant. Reaction (2) proceeds very rapidfy as can be seen from the corresponding forward reaction rate constant (k,) and reverse reaction rate constant (k_,) measured by Eigen ef 01. (1961) using a relaxation technique k, = 3.4 x 106s-’ k_, = 2.0 x 10s 1.mol-’ s-l
(20°C)
The rate constants of Wang and Him~lb~u (1964) seem to be too low by co 8 orders of magnitude 231
S
I
Table
Equlllbrmm constants ---__-_
GKAVE\HORST
and G
BLILKL
lor the system SO,-water -______._ __
as a lunctton
of temperature
Temperature
t CI
tmoli
5 15 25 Author’s
‘iatm-‘t 2 72
278 x lo-: 7 19 x 1o-2 174 x 10-L
I 76 I 2-l Gmehn
(1963)
Srllen
= 3.3 x IO-2s-‘; k-, = 2.71. mol-’ s-‘---xe for example Beilke and Lamb. 1975). Reactlon (3) is a pure ionic reaction and is therefore expected to proceed similarly fast as reaction (2). In Table 1 some values for the thermodynamic constants H. K, and K, are given as a function of temperature. Figure 1 shows the mole fractions of the three sulfur (IV) species m equilibrium as a function of solutlon pH at 25°C. As can be seen, the bisulfite ion HSO; is the most abundant sulfur (IV) species in the pH range of atmospheric droplets (i.e. pH 3-6). Transport of SOz within the gas phase to the drop let proceeds by turbulent and molecular difision. The rate determining molecular diffusion can be calculated using the Fickian equation for diffusion along with suitable initial and boundary conditions The SO*-concentration profile normal to the droplet surface for maximum diffusion can be calculated assuming zero concentration at the droplet surface (i.e. the droplet is a complete SO,-absorber) and constant S02-concentration in infinity: (k,
S(K) ----= ?I
(4)
?RZ
(5)
RC,fort=OandR>r
Boundary condition : RC=Ofort>OandR=r (6) where R = radius vector: I’ = droplet radius; C = S02-concentration in the gas phase; D = molecular diffusivity of SO, in air.* * D,,.
’
\sas tahen as O.lOcm’
b-
FIN. 1.
In our calculallons
of sulfur
IO c
Mole
fraction
100 )I 10.” 7 90 . lo- ” 624 * 10-B Sullen (1964)
(I 964)
This problem has been sohed by Smoluchowskl (1918) for diffusion of aerosol particles into a water droplet and has been applied to SO2 by Beilke (1965). The SO,-concentration profiles normal to the droplet are given by
c=c,
i-L+i R
x = Integration
(R-r),?,
2r R,;o
6
e-“‘dx
I
(IV) species
for
(7)
>
variable
and are shown in Fig. 2 for a droplet with a radius of IOprn after different times. The quantity of diffusing SO2, M,(g), which has entered the droplet in time r under the assumed conditions (4) to (6). is gven by: M,(g) = 4&C+
+ 5).
(8)
The assumption
D F’(RC) -
Initial condition: RC=
h, (molar)
‘I (molar)
that a water surface is a complete SO*-absorber (i.e. that the gas concentration is zero at the surface) is approximately valid for the absorption of SO* in alkaline solutions like ocean water (Liss, 1971; Liss and Slater, 1974; Brimblecombe and Spedding. 1972; Spedding, 1972; Beilke and Lamb. 1974). For atmospheric droplets with pH values normally between 3 and 6 with a limited absorption capacity for sulfur (IV), the droplet surface can represent a resistance to X&-transfer which is not negligible especially if the droplet pH is low (Liss, 1971; Brimblecombe and Spedding. 1972). If it is taken into account that the concentration of SO, increases at the droplet
m equilibnum PH.
at WC
as a function
of aqueous
solution
Heterogeneous SO,-oxidation
1
10
20
233
in the droplet phase
LO
30
R-r
Fig. 2. SO,-concentration profiles normal to the surface of a droplet (radius 10 q) diffusion calculated for a complete S02-absorber.
surface due to absorbed SO2 the calculated diffusion fluxes are reduced In Fig. 3 the SO,-flux to a complete absorbing droplet is compared with the SO,-flux to a droplet with increasing SO,-concentration at the surface according to the amount of SO2 absorbed It was assumed that the absorbed SO2 was immediately mixed within the droplet. In this figure the ratios between the SO,-fluxes to the droplet and the equilibrium concentration of s(W) are plotted as a function of time for various pH values which were kept constant during the absorption process In a droplet with r = 10pm equilibrium between SO2 in the gasphase and sulfur (IV) is reached within a few seconds Transport of SO2 within a droplet proceeds by molecular diffusion and turbulent transfer. If transport proceeds within the whole droplet by molecular diffusion only and if the SO,concentration at the surface is kept constant, the total amount of diffusing sulfur (IV) from the surface to the interior is given by
M40, 1 - f “
(9)
M,
* D was taken as 1.15 x IO-‘cm’s_’ in our calculation for 1O~C after Broeker and Peng (1974) for all sulfur (IV) species. t M , can be calculated on the basis of Equations (I J-(3).
10-‘
10“
SOJA
caused by molecular
(Crank, 1976) where r is droplet radius, D is molecular diffusivity of S(W) in water,* M,(a) is the quantity of diffusing S(IV) which has entered the droplet in time t, M,f is the concentration of S(N) in equilibrium with SO,-gasphase concentration. As seen in Fig. 4, for cloud and fog droplets smaller than 50~ radius, the equilibrium between SO2 in the gasphase and S(N) within the droplet is reached within approximately 1 s if the SO,concentration at the surface is kept constant. For cloud drops both the transport of SO2 to a droplet and transport of S(W) within a droplet will be enhanced by turbulent transfer. A comparison of calculated and measured ratios (Barrie. 1975) M,(a)/M, for a droplet with r = 1.06 mm radius shows a good agreement. On the basis of our estimations of SO,-transport both in the gaseous- and droplet phase equilibrium between SO, in the gas phase and sulfur (IV) in the droplet phase will be achieved within less than co I s for atmospheric cloud and fog droplets. As will be shown in the next section, oxidation of SJIV) to sulfate in a droplet proceeds slower by orders of magnitude. Thus the rate determining step in the overall heterogeneous SO,-oxidation in cloud and fog droplets smaller than 5Opm is the oxidation of S (IV) to sulfate and not diffusion processes.
10“ time
1
10
set
S. BEILKE
and G
GRAVFNHORST
or Rand and Gale (1967) but is m contrast tsl tha: of Schroeter (1963) and Penkett et al. (1976).
L
I
0.05
01
0.15
set
0.2
time
Fig. 4. Rattos between the quantrty of diffusing S(lV). M,(a). which has entered a droplet and quantrty of S(N) in equilibrium, MI. as a function of time and droplet radtus 3. OXIDATION
OF SO, IN THE DROPLET
(HSG;) or the sulfite ion (SO:-) is the oxygen carrier or both. A great many discrepancies in the literature concerning the pH dependance of SO1-oxidation and sulfate formation could be removed if there were a clear answer to this question since in a drop in equilibrium with a fixed SO2 concentration HSO; varies with [H’]-’ and SOi- with [H+]-‘. We will interpret the SO+xidation in terms of the more widely accepted view that the sulfite ion SO: - is the oxygen carrier and not HSO; or both. This assumption is in accordance for example with Winkelmann (1955) * krO is defined by: d[SO:-]/dt = klo [SO:-]. t In order to compare the results of Brimblecombe and their rate
pH 4 to 6
into our concept. $ For Penkett et 111.(1976) we calculated the corresponding k,, value for pH 6. 8 For Schroeter pH range 7-8
(1963) we calculated
k,.
values for the
i
Fig. 5. Pseudo first order rate constant for k,,,, as a function of pH.
According to Winkelmann (1955). the oxidatron follows the simple overall reaction.
PHASE
One of the most important questions in Sot-oxidation in the droplet phase is whether the bisulfite ion
Spedding (1974) with the others we transformed constants given in Table 1 of their paper for
In the following sections we wtll compare three dtfferent mechanisms of SO,-oxtdatlon and assess theu importance for sulfate formatron m the droplet phase In order to compare mvestrgattons. we have transformed all rate expressrons-whenever possible -Into the corresponding expresston of our concept wrth the sulfite ion (SO:-) being the oxygen earner.
reaction
10,
Most authors agree that the oxidation IS zero order with respect to oxygen over a wade range of 02-concentrations (for example Titoff. 1903; Riccoboni et al., 1949; Winkelmann, 1955; Penkett et al.. 1976). The order with respect to suUite is unity in most investigations (for example Titoff, 1903; Winkelmann, 1955; Rand and Gale, 1967). The following overall pseudofirst-order rate constants, k, 0, for reaction (10) are available* : (a) klo = 1.7 x 10-3s-’ at pH 6.8 and cn 25°C. Scott and Hobbs (1967). from measurements of Van den Heuvel and Mason (1963) at room temperature. between pH 2-4 at 25°C. (b) k,o = 3 x lo-‘s-r Miller and de Pena (1972). (c) klo = 3.5 x IO-‘s-l at pH 7 and 25°C. Winkelmann (1955). (d) k,,,t = 3.7 x iOT3-0.6 x 10-3s-’ between pH 4-6 at 25°C with 0.6 x lo-’ at pH 6. Brimblecombe and Spedding (1974). (e) k,J = 0.5 x 10e3 s-’ at pH 6 and 20°C. Pen kett ef al. (1976). (0 k,oij = C(I 6 x 10-3-0.6 x 10-3s-’ between pH 7-8 at 25°C. Schroeter (1963) (g) k,o = 1.2 x 10m4 [H+]-“.‘6 between pH 3-6 at 25°C. (k in s-l, [H’] in mol/l.). Beilke er al. (1975). (h) k,o = 0.013 + 59 [H+]‘!* (pH below 6 at 25°C) (k in s-‘, [H’] in mol/l.). MC Kay (1971) from measurements of Fuller and Crist (1941). Figure 5 shows klo values as a function of pH for 25°C. As seen in this figure, the k,,-values used by McKay (1971) and deduced from measurements of Fuller and Crist (1941) are higher than all other values by more than two orders of magnitude. One of the reasons for these high values is that Fuller and Crist did not take into consideration the variation of pH during the reaction nor the resulting changes of the ionic composition of the reaction mixture, even though such changes of the mole fraction of HSO; and SO:- are clearly indicated in Fig. 1
Heterogeneous
3
L
5
6
SO,-oxidation
PH
7
Fig. 6. Sulfate formation rates in the droplet phase as a function of droplet pH for three different SO,-oxidation mechanisms The calculated rates apply to the following conditions: gasphase concentrations for SOI: 1 ppb. O3 :40 ppb, T = 10°C.
For this reason we will confine ourselves to the other k,c-values given in the pH range 3 to 7. Although there is an opposite pH dependance of k,, between Beilke et al. (1975) and Brimblecombe and Spedding (1974). the values differ by no more than a factor of ca 2 in the pH range between 5 and 6. For pH 6 the corresponding value of .Penkett et al. (1976) is also within this range. Therefore it seems reasonable to accept a k10 value of ca 10-3s-’ in the pH range 5 to 6 at 25°C. Although the oxidation can be described by reaction (lo), the exact mechanism seems still to be a matter of discussion. It is likely that the oxidation proceeds via two chains with the radicals HO2 and SO; as chain carriers (Schmittkunz, 1963). Two results of Beilke er ol. (1975) who used a heterogeneous system to study SO,-oxidation should be mentioned here: Sulfate formation in aqueous solution was first order with respect to SO2 in the gas phase and did not depend on temperature within experimental error. This somewhat surprising result may well have arisen from the compensating effects of decreasing SO,-solubility (see Table I) and increasing reaction rate constant with increasing temperature. On the basis of this result an activation energy for reaction (10) has been estimated to be 14 kcal mol-’ in the pH range between 3 and 6 using the thermodynamic quantities given in Table 1. Since Beilke et al. (1975) used a heterogeneous system. their laboratory results will be directly applied to the atmosphere. For a cloud with a liquid water content (LWC) of 0.1 g rne3, S02-removal rates were calculated to fall between ca 10-6% h-i and 1.5% h-l for droplet pH values between 3 and 6, or ten times higher for LWC = 1 g rnb3. The lower dashed curves in Fig 6 show sulfate formation rates due to SO,-oxidation by O2 in the absence of catalysts as a function of droplet pH for 10°C and 1 ppb SO,-gas phase concentration. This SO,-gas phase concentration was
in the droplet phase
235
found to be+realistic at cloud level (1-5 km) in background air (Georgii and Jo& 1964; Gravenhorst. 1975). As seen in this figure, sulfate formation rates are extremely small except for high droplet pH (above 5 6) In Fig. 6 sulfate formation rates which were calculated on the basis of the experimental data of Brimblecombe and Spedding (1974) and of Penkett et al. (1976) are also included. A comparison of the three dashed curves shows that very similar conclusions concerning the role of SO,-oxidation by O2 in the absence of catalysts can be drawn. A special type of the SO,-oxidation by 0, in the absence of catalysts is the oxidation in the presence of ammonia (NH,). In the past, this mechanism has attracted substantial attention. It had previously been accepted by many scientists that the so-called SO,NHS-H,O-mechanism was the most effective SOroxidation process in atmospheric droplets. The e&t of ammonia is to keep the droplet pH high when S(IV) is oxidized to sulfate. This effect was investigated by many scientists, for example by Junge and Ryan (1958), Van den Heuvel and Mason (19633 Scott and Hobbs (1967). MC Kay (1971). Miller and de Pena (1972) Easter and Hobbs (1974) and Tomasi et al. (1975). The model of Scott and Hobbs was in best agreement with measurements by Beilke et al. (1975) It can, however, not be used (as for example Easter and Hobbs (1974) did) in order to calculate sulfate formation rates without further modifications of the physics involved (Beilke and Barrie, 1974) The calculated sulfate formation rates of Easter and Hobbs apply to a droplet pH of 6.6 (Harrison, 1974). It seems to be unlikely that such a high pH value is generally found. In a real atmosphere NH3 competes for control over droplet pH with many other soluble substances for example with an acidic background aerosol (Junge and Scheich, 1969; Gravenhorst. 1975) causing in most cases an overall droplet pH between ca 4-5. The sulfate formation rates in the droplet phase should therefore be lower than those predicted in the model of Scott and Hobbs by a factor between co 10’ to 10’ (Beilke, to be published). Therefore the major function of NH3 is likely to be no longer an effective enhancement of the SO,-oxidation but rather the conversion of a pre-existing sulfate containing droplet into a droplet containing ammonium and sulfate ions. However. the absorption of ammonia can shift the pH value from the 3 to 4 range to a range where oxidation of SO1 by O2 in the presence of catalysts is more important (45). If the source of sulfate in cloud drops is aqueous phase oxidation the oxidation mechanism must be effective between pH 3 and 5. 3.2 SO,-oxidation by 0,
in the presence oj catalysts
Catalysts are mainly transition metals of the 4th period especially manganese and iron. In the case of catalytic S02-oxidation by O2 in the presence of cata-
S. BEILKEand G.
736
lysts we will clearly separate aqueous phase SO,-OXIdation occurring in saline solutions of deliquesced aerosol salts with catalyst concentration of lo- ‘-1O-3 M from that occurrmg in dilute solutions like fog and cloud drops with catalyst concentrations between 1O-4-1O-8 M. We will confine ourselves to the latter. Since the important work of Bigelow m 1897 more than 200 publications have appeared in the literature dealing with catalytic oxidation of S(N). Most worked with S02- and catalyst concentrations which were higher by orders of magnitude than those occurring m the atmosphere. In addition, most studies were carried out with sodium sulfite solutions which have a pH range between 8-9. The atmospheric pH-range is between 3 and 6. Therefore we will restrict discussion to a few investigations carried out with concentrations of SO2 and catalysts similar to atmospheric conditions (catalyst concentrations between 1O-4-1O-s M and (or) SO,-concentration far below 1 ppm). The effect of catalysts was mvestigated using both a heterogeneous system (Barrie and Georgii, 1976) and a homogeneous one (Coughanowr and Krause, 1965; Bracewell and Gall, 1967; Tsunogai, 1971; Brimblecombe and Spedding. 1974; Penkett et al.. 1975; Lenelle. 1975; Betz. 1977). The oxidation of S(N) was found to be proportional to the square of the Mn2+-concentration when the concentration ratios S(IV)/Mn*’ 9 1. (Coughanowr and Krause, 1965) which is in agreement with the findings of Bracewell and Gall (1967). An empirical equation describing the oxidation in the presence of Mn” was given by Bracewell and Gall: B 1 A z = @iiy
+ [s(w)];.2
where A, B = constants, [Mn’+]e = mitial concentration of Mn SO., in aqueous solution, [S(IV)Je = Initial S(IV) concentration in aqueous solution and R = reaction rate. Btimblecombe and Spedding (1972) investigated the Fe3’-catalyzed SO,-oxidation in acidic aqueous solutions with concentration for Fe’+ = 10e6 M and S(IV) 2 lo-’ M. They found the disappearance rate of S(N) to be proportioned to both Fe3’ and S(IV). At a pH of 4 the oxidation proceeded faster in the presence of 10m6 M Fe 3 ’ than in the absence of catalysts by ca a factor 103. Barrie and Georgii (1976) investigated the catalyzed SO*-oxidation at 25°C and 8°C using single droplets of 2.1 mm diameter conaisting of distilled water with heavy metal concentrations in the range between low6 to IO-” molar for manganese and iron. The droplets were exposed to SO+oncentrations in air between 0.01-l ppm. The sulfate * For a droplet with 1Opm radius the calculated rate of SO,-diffusion to the droplet IS higher than the measured SOI-oxidation rate by a factor between ca IO’ to IO4 for the droplet pH range of interest here (3-5).
GRAVENHORST
formation rate was found to be first order with respect to SO2 in the gas phase. In the pH range 24 5 the most effective catalyst was Mn2’ followed hk Fe” and to a much lesser extent Fe’ + The catalytic effectiveness of Mn” was enhanced by the addition of Fe’+ or Fe3’ This synerglsttc effect was also observed by other authors (Penkett r~ 01. 1975. Lenelle, 1975). The Increase of temperature of a solution
of Mn’+ from 8 to 25°C caused an Increase in OXIdation rate by a factor 5 to IO m the pH range 2 4.5. However, the temperature dependance was reduced when Fe*’ or Fe3’-ions were added to the Mn’+ solution. On the basis of the experiments of Barrie and Georgia (1976), the conclusion can be drawn that sulfate formation due to catalytic SO,-oxidation can be significant in urban air clouds or fogs in which case the catalysts concentration can be of the order of 10m5 molar or even higher. A realistic approach with respect to concentrations and composition of catalysts is to use natural rainwater as solvent for studying the catalytic S02-oxidation. Betz (1977) has investigated the oxidation in a homogeneous aqueous phase of rainwater collected in the area of Frankfurt/M. (Germany). The heavy metal concentrations were between 10-7-l0-b molar for manganese and between 1O-6-1O-5 molar for iron. Of the total manganese SO-90% were &ssolved. For iron the corresponding fraction was 60-75%. The rainwater pH was between 3.2 and 5.2. By measuring the concentration decrease of S(N), sulfate formation rates were determined as a function of pH and temperature. The activation energy was measured to be 23 kcal mol- ’ which is in close agreement with the value of 21 kcal mol- ’ found by Penkett et cl. (1975). The rate constants are ca 1 to 2 orders of magnitude higher than those reported for the SO,-oxidation by O2 in the absence of metal catalysts. In order to apply the experimental data of Betz (1977) into the atmospheric heterogeneous system. three assumptions were made: (1) SOz-transport processes proceed fast compared with oxidation in the droplet phase.* (2) Sulfate formation proceeds first order with respect to SO2 in the gasphase. (3) The concentrations of heavy metals measured in rainwater are the same as in cloud or fog water. This assumption is to some extent uncertain. The results of our calculations (solid curve m Fig. 6) for a gas phase concentration of 1 ppb and 10°C show that sulfate formation proceeds faster than in the absence of metal catalysts by CO2 orders of magnitude. It has often been concluded on the basis of experiments with natural rainwater that this type of SO1-oxidation contributes only little to sulfate formation in the droplet phase at least for country clouds in rural areas. These conclusions are drawn under the assumption that catalyst concentrations and pH in rainwater are the same as in cloudwater which is highly uncertain. Very similar results have been
Heterogeneous
SO,-oxidation
obtained by Penkett et af. (1975a) and Lenelle (1975).* The oxidation rate measured by Tsunogai (1971) seems to be lower. The mechanism of the catalyzed %&oxidation by O2 is still a matter of discussion. The oxidation proceeds likely via two chains with the radicals HO* and SO, as chain carriers (Schmittkunz, 1963). The radicals are likely produced as follows: (see Barrie and Georgii (1976) who used the proposed m~h~isrn of ~hrni~k~ (1963)).
[MeX+ S(IV),J-‘6-“’ + 0,-r [Mer+ S(IV)~]-‘5-X’ + 0, 0;
+ H+-+HO*
[Me”’ S(IV)~]-‘5-X’ + SOi- - [Me”’ S(W),]-(“-*)
(14)
As can be seen, the complex [Me”’ s(IV),]-‘6-x) is regenerated Me ++ is transition metal ion of valency x. S(IV) is SO:- or HSO;. In the case of manganese as a catalyst, the results of Barrie and Georgii (1976) Seem to indicate that the manganese ion Mn 2+ formed complexes [Mn2+ (SO;-)s]-4 before taking part in the reaction. Interested readers are referred to the publication of Schmittkunz (1963) 3.3 SOL-ox~uriun by strongly oxidizing agem The effect of strongly oxidizing agents was investigated by Penkett (1972); Penkett and Garland (1974); Penkett, Jones and Eggleton (1975); Penkett, Brice and Eggleton (1976) and of Barrie (1975) for ozone (0,) and by Penkett, Brice and Eggleton (1976) for hydrogen peroxide (H202j The reaction of ozone with sodium sulfite in a homogen~us aqueous phase was studied by Penkett (1972) at pH 4.65 and was extended later to a pH range 1 to 5 by Penkett, Jones and Eggleton (1975). The SO,-oxidation was measured by following the decline in ozone. In these experiments the concentration of total sulfitet was kept in excess of ozone. In this case the reaction was first order with respect to ozone. The order with respect to total sulfite was unity at pH 4 and decreased with increasing PH. The O,-concentrations used were about 5 x tom5 M cor-
* Since the papers of Penkett er al. (1975) and of Lenelle (1975) are not yet published in the open literature, no further details on their work are gMn here. t Total sulfite is mainly HSO; in the range investigated by Penkett et al. @H 1-S). $ For a droplet with r = fO#rn radius the calculated rate of SOrdifFusion to the droplet is higher than the measured (Penkett et al., 1975) SO,-oxidation rate due to ozone by a factor between ca 10’ to 10’ for the droplet pH range 3-5. P Since the results of Pen&t, Brice and Eggleton (1976) have not been published yet in the open literature, no details are given here.
237
responding to a gas phase concentration of 4 2700 ppm in equilibrium. According to Penkett and Garland (1974) the results obtained with such high Osconcentrations can be extrapolated to atmospheric Os~onc~trations as indicated by their measurements on the rate of SO,-oxidation in a fog chamber using atmospheric concentrations of both ozone and SO*. The mechanism of the reaction between ozone and total sutfite is likely to proceed via a free radical mechanism (Penkett, Jones and Eggleton
(12) (13)
+ SO;
in the droplet phase
1975).*
It should be mentioned, however, that Barrie (1975) did not find a much enhanced SO+xidation rate due to ozone at a pH of co 4. He exposed three droplets (1.06mm radius) of distilled water to a stream of air containing a mixture of 0.4ppm SO, and different Os-concentrations between 0.01 and 0.5 ppm. In the droplet phase the increase of total sulfur was measured using a new isotope dilution technique (Klockow et al.. 1974). Barrie’s measurements with distilled water droplets showed only a very low enhancement of SO,-oxidation rate due to ozone. He found, however, an enhancement of the oxidation rate due to O5 when manganese was present. For an O,-gasphase concentration of 0.14 ppm the rate in a 10W5molar MnCl,droplet was doubled In order to estimate sulfate formation rates for the heterogeneous atmospheric droplet system on the basis of the experiments of Penkett (1972) and Penkett, Jones and Eggleton (1975) using a homogeneous aqueous phase. we used their kinetic data along with the thermodynamic quantities in Table 1 for S(W) and for ozone (Barrie, 1975j Again, we assumed transport processes for both SO1-gas and OS-gas to proceed fast compared with the SO,-oxidation1 in the droplet The uppermost dashed curve in Fig. 6 shows sulfate fo~tion rates due to SO,-oxidation by ozone after the experimental results of Penkett, Jones and Eggleton (1975) for the following set of conditions: gas phase concentrations are 1 ppb for SO, and 40 ppb for ozone and 10°C. The corresponding sulfate formation rate calculated on the basis of Barrie’s experiment is lower by a factor of co lo* at pH 4. Another oxidizing agent of importance for S02-oxidation in the droplet phase could be H20t for two reasons: (a) The solubility of H102 in water is extremely high. (b) The reported steady state concentrations of H20, in the lower troposphere are about 1 ppb (3 ppb according to model calculations of Levy (1973) and 1 ppb after Crutzen (1977) (persona1 communication to S. Beilke). Very recent experiments of Penkett, Brice and Eggleton (1976% have shown that the reaction is very fast even in the pH range 3 to 5. However, the importance of this oxidation process for sulfate formation in the droplet phase depends strongly on H202 gasphase con~~a~ons which have not yet been measured to our knowledge.
23X
5 BEILK~ and Ci. GRAVENHORST 4. IMPORTANCE SOI-OXIDATION
OF HFZTEROCENEOLS
IN THE ATMOSPHERE
FURTHER
AND
RESEARCH
Figure 6 shows a comparison of sulfate formation rates due to three S02-oxidation mechanisms (a) SO,-oxidation by O2 m the absence of catalysts (lower dashed curves). (b) S02-oxidation by O2 in the presence of catalysts (solid curve) and (c) S02-oxidation by ozone (upper dashed curve and one single dot) As seen in Fig. 6. oxidation type (a) IS unimportant for sulfate formation unless droplet pH is higher than ca 6. Although sulfate formation by type (b) proceeds faster by two orders of magnitude than bj type (a). it seems to be at least for background conditions of little importance for atmospheric sulfate formation unless droplet pH is higher than cu 5. For urban fog conditions, however. with higher heavy metal concentrations. this mechanism may play an important role. Sulfate formation rates due to S02-oxidation by ozone (type (c)) are higher than the corresponding rates for type (b) by ca 1 order of magnitude accordmg to Penkett er a/. (1975). In contrast. however, Barrle (1975) found that at least at pH 4 this rate is lower. The main discrepancies in heterogeneous SO,oxidation obvious from Fig. 6 are the pH dependance and the role of ozone. Therefore we suggest the following further research. (a) It should be determined whether the SOi--ion or the HSOJ-ion (or both) carries the oxygen in the pH range 3 to 7 This question is of fundamental Importance for the pH dependance of sulfate formation m atmospheric droplets. (b) Further research is necessary to explain the dlscrepancy of the experimental results of Penkett et al. (1975) and Barne (1975) concerning the role of ozone m heterogeneous SO,-oxidation in the absence and presence of catalysts. (c) In order to calculate sulfate formation rates due to SOI-oxidation by O2 in the presence of catalysts, we assumed that the concentrations of heavy metals and the pH value measured in rainwater are the same as m cloudwater. We therefore recommend measurements of the composltlon and pH of cloud and fog water since the small cloud and fog droplets are more important for sulfate formation due to their relatively long lifetime (minutes to hours) compared with the lifetime of large ramdrops (ca I minute). In a real atmosphere the three oxidation mechanisms proceed at the same time. We are therefore concerned with a very complex synergistic S02-oxidation by 02, 03. H,O, m the presence of NHJ, metal catalysts and organic compounds. Organic substances may either act as inhibitors or catalyst for SO,-oxidation and can increase or de* Droplet pH can easdy be measured with a micro-pH cell t Total sulfur m a smgle droplet can be measured with the Isotope dllutlon technique of Klockow YI al (1974)
crease droplet pH. Any ammonia absorbed by a droplet tends to increase the droplet pH which m turn causes higher sulfate formation rates due to oxidation by O2 in the presence of catalysts since the amount of S(IV) increases for a given SO,-gas phase concentration. In addition. the catalytic effectivity of any heavy metal of the 4th period depends also on pH. Furthermore. the catalytic effectivity of a special metal ion depends on the presence of other metals. The SO,-oxidation
rate obtained
with a mixture
of
iron IS. for example, higher than the sum of the rates with each ion present alone (Barne. 1975). In addition to the influence of pH the promoting effect of ozone for S02-oxidation, or the mhibiting effect of NO2 both m the presence of metal catalysts (Barrie. 1975) complicates the situation. (d) Because the overall effect of SO,-oxidation m droplets depends on the complex variety of interacbans, experiments should be carned out which take all these possibilities mto account. manganese and
These experiments should complement the studies suggested under (aHc) above. Such experiments should be carried out in a heterogeneous system. since m a homogeneous aqueous system compounds reacting in the liquid phase cannot be supplied from the gasphase to replace removed ones. That IS for example true for SOz. 03. H,02 and NH3. We suggest that drops consisting of natural rain- or cloudwater be suspended in outside air to simulate actual atmospheric conditions. By measuring the change of droplet pH* and increase of total sulfurt or sulfate, sulfate formation rates can be determined under realistic conditions. REFERENCES
Barrle L. (1975) An experimental investigation of the absorption of sulfur dioxide by cloud- and rain drops containing heavy metals. Berichte des Institutes fiir Meteorologie der Universitiit Frankfurt a.M.. Nr. 28. Barrle L. and Georgii H.-W. (1976) An experimental investigation of the absorption of sulfur dioxide by water drops containing heavy metal ions. Atmospheric Enuironment 10, 743-749. Beilke S. and Barrie L. (1974) On the role of NH3 m heterogeneous SO,-oxidation in the atmosphere. Paper presented at the European sulfur symposium, Ispra. Italy, October 1974. Beilke S. (1965) Untersuchungen liber ‘rainout’ und ‘washout’ von Schwefel-dioxid. Diplomarbeit. UniversitHt Frankfurt a.M. Beilke S. SO1-oxidation m the aqueous phase. To appear as a book m late 1978 Be&e S. and Lamb D. (1974) On the absorption of SO1 in ocean water. Tellus XXVI. 268-271 Beilke S. and Lamb D. (197Sj Remarks on the rate of formation of bisulfite ions in aqueous systems. AIChE Journal 21, 402404. Beilke S.. Lamb D. and Miiller J. (1975) On the uncatalyzed oxidation of atmospheric SO2 by oxygen in aaueous svstems. Atmosdwric Enoironment 9, 1083-1090. Be& M. (1957) Untersuchbngen iiber die Absorption und Oxidation von Schwefeldioxid in nattirlichem Regenwasser. Diplomarbeit. Universitlt Frankfurt September 1976
Heterogeneous
SO,-oxidation
Bracewell J. M. and Gall D. (1967) The catalytic oxidation of sulphur dioxide in solution at concentrations occurring in fog droplets Proc. Symp. on Physicochemical Transformation of Sulphur Dioxide Compounds in the Atmosphere and the Formation of Acid Smogs. Mainz. Germany, 1967. Brimblecombe P. and Spedding D. I. (1972) Rate of solution of gaseous sulphur dioxide of atmospheric concentrations. Nature 236, 225. Brimblecombe P. and Spedding D. 1. (1974) The catalyzed oxidation of micromolar aqueous sulphur dioxide-I. Atmospheric
Environment
8, 937-945.
Broecker W. S. and Peng T. H. (1974) Gas exchange rates between air and sea. Tel/us XXVI, 21-35. Coughanowr D. R. and Krause F. E. (1965) The reaction of SO2 and O2 in aqueous solutions of Mn S04. Ind. Eng. Chem. Fund. 4, 61-61.
Crank I. (1976) The Mathematics of &&ion. 2nd edn. Clarendon Press, Oxford. Crutzen P. (1977) Personal communication to S. Beilke. Easter R. C. and Hobbs P. (1974) The formation of sulphates and the enhancement of cloud condensation nuclei in clouds. J. atm Sci. 31, 1586-1594. Eigen M., Kustin K. and Maass G. (1961) Die Geschwindigkeit der Hydratation von SOI in wPssriger LSsung. Z. Phys. Chem 30, 130. Fuller E. C. and Crist R. H. (1941) The rate of oxidation of sulfite ions by oxygen. J. Am. them. Sot. 63, lWl650.
Georgii H. W. and Jost D. (1964) Untersuchungen iiber von Spurengasen in der freien die Vertklung Atmosphire. Pure appl. Geophys 59, 217-224. Gravenhorst G. (1975) Der Sulfatanteil im atmosphfr&hen Aerosol iiber dem Nordatlantik. Berichte des Institutes riir Meteorologie der Universitiit Frankfurt a.M. Nr. 29 Gmelin (1963) Handbuch Schwefel, 8. Auflage, Teil B-Lieferung 3. Harrison H. (1974) Discussion to paper by S. Beilke, D. Lamb, and J. Miiller, The heterogeneous oxidation of SO, in relation to atmospheric scavenging. Precipitation Scavenging. U.S. Atomic Energy Commission, Oake Ridge, Tenn. ERDA symposium series. CONF-741003. Junge C. E. and Ryan T. G. (1958) Study of the S02-oxidation in solution and its role in atmospheric chemistry. Q. Jl. R. met. Sot. 84, 4655. Junge C. E. and Scheich G. (I 969) Studien zur Bestimmung des S;iuregehaltes von Aerosolteilchen Atmospheric Environment
3, 423441.
Klockow D., Denzinger H. and Riinicke R. (1974) Anwendung der substtihiometrischen Isotopenverdiinnungsanalyse auf die Bestimmung von atmosphlrischem Schwefel und Chlor in Background Luft Chem. Ing. Techn. 46. 18. Lenelle Y. (1975) Contribution a I’etude de l’oxydation des sulfites dans la pluie. Paper presented at the European sulfur symposium. Ispra, Italy, October 1975. Levy H. (1973) Photochemistry of minor constituents in the troposphere. Planet. Space Sci. 21, 575-591. Liss P. S. (1971) Exchange of SO2 between the atmosphere and natural waters Nature 233, 327-329. Liss P. S. and Slater P. G. (1974) Flux of gases across the air-sea interface. Natwe 247, 181-184. McKay H. A. C. (1971) The atmospheric oxidation of sul-
239
in the droplet phase
phur dioxide in water droplets in presence of ammonia. Atmospheric
Environment
!!, 7-14.
Miller J. M. and de Pena R. G. (1972) Contribution of scavenged sulfur dioxide to the sulfate content of rain water. J. geophys. Res. 77. 5905-5916. Penkett S. A. (1972) Oxidation of SO, and other atmospheric gases by ozone in aqueous solution. Nature Phys. Sci. 240. 105-106. Penkett S. A. and Garland J. A. (1974) Oxidation of sulphur dioxide m artificial fogs by ozone. Tel/us XXVI, 284-290. Penkett S. A., JonesB. M. R. and Eggleton A. E. J. (1975a) A study of SO,-oxidation in stored rainwater samples Paper presented at the European sulfur symposium, Paris, 1975. Penkett S. A., Jones B. M. R. and Eggleton A. E. J. (1975b) Rate of oxidation of sodium sulphite solutions by oxygen and by ozone. Paper presented at the European sulfur symposium, Ispra, 1975. Penkett S. A., Brice K. A. and Eggleton A. E. J. (1976) A study of the rate of oxidation of sodium sulphite solution by hydrogen peroxide and its importance to the formation of sulphate in cloud- and rainwater. Paper presented at the European sulfur symposium, Ispra 1976. Rand M. C. and Gale S. B. (1967) Kinetics of the oxidation of sulfites by dissolved oxygen. In Principles and Applications of Water Chemistry (Edited by Faust S. D. and Hunter J. V.) pp. 380404. Wiley, New York. Riccoboni L., FotIani A. and Vecchi E. (1949) Studi di cinetica chimica sui processi di autoossidazione in fase liquida diluita-II. Gazz. chim. Ital. 79, 418-442. Schmittkunz H. (1963) Chemlumineszenz der Sulfitdxidation. Dissertation of the Naturwissenschaftliche Fakultit der Universitit Frankfurt. Schroeter L. C. (1963) Kinetics of air oxidation of sulfurous acid salts. .I. Pharm. Sci. St, 559-563. Scott W. D. and Hobbs P. V. (1967) The formation of sulfate in water droplets 1. atm. Sci. 24, 5457. Sillen L. G. (1964) Stability Constants of Metal-Ion Complexes. Section I-inorganic Ligands, second edition. Chem. Sot. (London), Special Publication No. 17. Smoluchowski M. (1918) Versuch einer mathematischen Theorie der Koagulationskinetik. Z. Phys. Chem. 151. Spedding D. J. (1972) Sulfur dioxide absorption by sea water. Atmospheric Enoironment 6, 583-586. Titoff A. (I 903) Beitriige zur Kenntnis der negativen Katalyse in homogenen System. Z. Phys. Chem. 45, 641-683. Tomasl C., Guui R. and Vittorr 0. (1975) The ‘SOI-NH,solution droplet’ system in an urban atmosphere. J. arm. Sci. 32, 1580-1586.
Tsunogai S. (1971) Oxidation of sulfite in water and its bearing on the origin of sulphate in meteoric precipitation. Geochem. J. 5, 175. van den Heuvel A. P. and Mason B. J. (1963) The formation of ammonium sulfate in water droplets exposed to gaseous sulfur dioxide and ammonia. Q. JI. R. met. SOC. 89, 271-275. Wang J. C. and Himmelblau D. M. (1964) A kinetic study of sulfur dioxide in aqueous solution with radioactive tracers. AIChE Journal 21.402-404. Winkehnann D. (1955) Die elektrochemische Messung der Oxidationsgeschwindigkeit von Na2S0, durch geliisten Sauerstoff. Z. Eleklrochemie 59. 891-895.
IN THE
S. BEILKE Umweltbundesamt:
Pilotstation
Frankfurt. FeldbergstraDe 45, 6000 Frankfurt a.M.. W. Germany
and G. GRAVENHORX Institut fiir Atmosphiirische Chemie, Kemforschungsanlage Jiilich, W. Germany Abstract-In order to determine the rate controiling step for sulfate formation in a heterogeneous droplet system. Sot-transfer within the gasphase towards and within droplets are calculated. equilibrium between SO2 in the gasphase and sulfur (IV) in cloud and fog droplets is reached within less than 1 second. Oxidation of sulfur (IV) to sulfate in the droplet phase proceeds slower by orders of magnitutde. Three mechanisms of SO,-oxidation are discussed: {a) SO,-oxidation by Ot in the absence of catalysts; (b) SO,-oxidation by O2 in the presence of catalysts, and (cc) SO,-oxidation by strongly oxidizing agents. Mechanism (a) contributes only to a negligible extent to sulfate formation in droplets even in the presence of typical concentrations of ammonia. Indications are strong that the major function of the SO,-NH,-&O-system is not the oxidation of SO2 to sulfate. Oxidation mechanism (b) may contribute to a significant extent to sulfate formation in urban fogs in which case the concentrations of catalysts can be sufficiently high. For clouds in remote areas with much lower catalyst concentrations SO,-oxidation by mechanism (b) seems to be of little importance. The oxidation by strongly oxidizing agents (mechanism (c)) appears to be the dominant mechanism although some experimental discrepancies have to be resolved.
1. lNTRODtJCllON There is strong evidence that the precinct part of atmospheric sulfate found in the Federai Republic of Germany is generated by oxidation of atmospheric SO, and not directly emitted as sulfate containing particles. The oxidation proceeds by 2 processes: (I) Homogeneous SO,-gasphase oxidation. (2) Heterogeneous SOI-oxidation in atmospheric droplets and on aerosol particles This paper deals with the SOIoxidation in the heterogeneous SO,-air-droplet-systern. It is attempted to integrate some of the reported experimental and theoretical investigations into a unified picture and to assess some quantitative data on sulfate formation and SO,-removal in the atmosphere. In a real atmosphere we are concerned with a heterogeneous SOI-oxidation in the aqueous phase whereas in laboratory experiments both homogeneous and heterogeneous aqueous systems were used to study SO,-oxidation. In order 10 apply the laboratory results obtained in a homogeneous aqueous phase to the atmospheric heterogeneous system some estimates are made of SO,-transport rates within both the gaseous and the droplet phase. These calculations are described in the first part of this paper. In the second part. three mechanisms of SOz-oxidation will be discussed and their importance for atmospheric sulfate formation assessed: (a) SO,-oxidation * Sulfur (IV) stands for the sum of SO,-H20, and SO:-.
HSO;
by O2 in the absence of catalysts, (b) SO1-oxidation by O2 in the presence of catalysts and (c) Sot-oxidation by strongly oxydizing agents. On the basis of this discussion further studies to 611 knowledge gaps will be suggested. 2 SO,-TRANSFER AND ARSORPTION
In a heterogeneous system in which gaseous SO2 is in equilibrium with sulfur (IV)* in the aqueous phase, the following equations describe the equilib rium, (SO,),,, + Hz0 &
SOz+I,O
SO~+z*o AH+ 7 HSO; KI,H+
(1)
-f HSO;
(2)
+ SO:-
(3)
where SOz*H20 is physically-halved SOz, HSO; is the bisulfite ion, and SOi- is the sulfite ion. H = Henry constant; K, = first dissociation constant; K, = second dissociation constant. Reaction (2) proceeds very rapidfy as can be seen from the corresponding forward reaction rate constant (k,) and reverse reaction rate constant (k_,) measured by Eigen ef 01. (1961) using a relaxation technique k, = 3.4 x 106s-’ k_, = 2.0 x 10s 1.mol-’ s-l
(20°C)
The rate constants of Wang and Him~lb~u (1964) seem to be too low by co 8 orders of magnitude 231
S
I
Table
Equlllbrmm constants ---__-_
GKAVE\HORST
and G
BLILKL
lor the system SO,-water -______._ __
as a lunctton
of temperature
Temperature
t CI
tmoli
5 15 25 Author’s
‘iatm-‘t 2 72
278 x lo-: 7 19 x 1o-2 174 x 10-L
I 76 I 2-l Gmehn
(1963)
Srllen
= 3.3 x IO-2s-‘; k-, = 2.71. mol-’ s-‘---xe for example Beilke and Lamb. 1975). Reactlon (3) is a pure ionic reaction and is therefore expected to proceed similarly fast as reaction (2). In Table 1 some values for the thermodynamic constants H. K, and K, are given as a function of temperature. Figure 1 shows the mole fractions of the three sulfur (IV) species m equilibrium as a function of solutlon pH at 25°C. As can be seen, the bisulfite ion HSO; is the most abundant sulfur (IV) species in the pH range of atmospheric droplets (i.e. pH 3-6). Transport of SOz within the gas phase to the drop let proceeds by turbulent and molecular difision. The rate determining molecular diffusion can be calculated using the Fickian equation for diffusion along with suitable initial and boundary conditions The SO*-concentration profile normal to the droplet surface for maximum diffusion can be calculated assuming zero concentration at the droplet surface (i.e. the droplet is a complete SO,-absorber) and constant S02-concentration in infinity: (k,
S(K) ----= ?I
(4)
?RZ
(5)
RC,fort=OandR>r
Boundary condition : RC=Ofort>OandR=r (6) where R = radius vector: I’ = droplet radius; C = S02-concentration in the gas phase; D = molecular diffusivity of SO, in air.* * D,,.
’
\sas tahen as O.lOcm’
b-
FIN. 1.
In our calculallons
of sulfur
IO c
Mole
fraction
100 )I 10.” 7 90 . lo- ” 624 * 10-B Sullen (1964)
(I 964)
This problem has been sohed by Smoluchowskl (1918) for diffusion of aerosol particles into a water droplet and has been applied to SO2 by Beilke (1965). The SO,-concentration profiles normal to the droplet are given by
c=c,
i-L+i R
x = Integration
(R-r),?,
2r R,;o
6
e-“‘dx
I
(IV) species
for
(7)
>
variable
and are shown in Fig. 2 for a droplet with a radius of IOprn after different times. The quantity of diffusing SO2, M,(g), which has entered the droplet in time r under the assumed conditions (4) to (6). is gven by: M,(g) = 4&C+
+ 5).
(8)
The assumption
D F’(RC) -
Initial condition: RC=
h, (molar)
‘I (molar)
that a water surface is a complete SO*-absorber (i.e. that the gas concentration is zero at the surface) is approximately valid for the absorption of SO* in alkaline solutions like ocean water (Liss, 1971; Liss and Slater, 1974; Brimblecombe and Spedding. 1972; Spedding, 1972; Beilke and Lamb. 1974). For atmospheric droplets with pH values normally between 3 and 6 with a limited absorption capacity for sulfur (IV), the droplet surface can represent a resistance to X&-transfer which is not negligible especially if the droplet pH is low (Liss, 1971; Brimblecombe and Spedding. 1972). If it is taken into account that the concentration of SO, increases at the droplet
m equilibnum PH.
at WC
as a function
of aqueous
solution
Heterogeneous SO,-oxidation
1
10
20
233
in the droplet phase
LO
30
R-r
Fig. 2. SO,-concentration profiles normal to the surface of a droplet (radius 10 q) diffusion calculated for a complete S02-absorber.
surface due to absorbed SO2 the calculated diffusion fluxes are reduced In Fig. 3 the SO,-flux to a complete absorbing droplet is compared with the SO,-flux to a droplet with increasing SO,-concentration at the surface according to the amount of SO2 absorbed It was assumed that the absorbed SO2 was immediately mixed within the droplet. In this figure the ratios between the SO,-fluxes to the droplet and the equilibrium concentration of s(W) are plotted as a function of time for various pH values which were kept constant during the absorption process In a droplet with r = 10pm equilibrium between SO2 in the gasphase and sulfur (IV) is reached within a few seconds Transport of SO2 within a droplet proceeds by molecular diffusion and turbulent transfer. If transport proceeds within the whole droplet by molecular diffusion only and if the SO,concentration at the surface is kept constant, the total amount of diffusing sulfur (IV) from the surface to the interior is given by
M40, 1 - f “
(9)
M,
* D was taken as 1.15 x IO-‘cm’s_’ in our calculation for 1O~C after Broeker and Peng (1974) for all sulfur (IV) species. t M , can be calculated on the basis of Equations (I J-(3).
10-‘
10“
SOJA
caused by molecular
(Crank, 1976) where r is droplet radius, D is molecular diffusivity of S(W) in water,* M,(a) is the quantity of diffusing S(IV) which has entered the droplet in time t, M,f is the concentration of S(N) in equilibrium with SO,-gasphase concentration. As seen in Fig. 4, for cloud and fog droplets smaller than 50~ radius, the equilibrium between SO2 in the gasphase and S(N) within the droplet is reached within approximately 1 s if the SO,concentration at the surface is kept constant. For cloud drops both the transport of SO2 to a droplet and transport of S(W) within a droplet will be enhanced by turbulent transfer. A comparison of calculated and measured ratios (Barrie. 1975) M,(a)/M, for a droplet with r = 1.06 mm radius shows a good agreement. On the basis of our estimations of SO,-transport both in the gaseous- and droplet phase equilibrium between SO, in the gas phase and sulfur (IV) in the droplet phase will be achieved within less than co I s for atmospheric cloud and fog droplets. As will be shown in the next section, oxidation of SJIV) to sulfate in a droplet proceeds slower by orders of magnitude. Thus the rate determining step in the overall heterogeneous SO,-oxidation in cloud and fog droplets smaller than 5Opm is the oxidation of S (IV) to sulfate and not diffusion processes.
10“ time
1
10
set
S. BEILKE
and G
GRAVFNHORST
or Rand and Gale (1967) but is m contrast tsl tha: of Schroeter (1963) and Penkett et al. (1976).
L
I
0.05
01
0.15
set
0.2
time
Fig. 4. Rattos between the quantrty of diffusing S(lV). M,(a). which has entered a droplet and quantrty of S(N) in equilibrium, MI. as a function of time and droplet radtus 3. OXIDATION
OF SO, IN THE DROPLET
(HSG;) or the sulfite ion (SO:-) is the oxygen carrier or both. A great many discrepancies in the literature concerning the pH dependance of SO1-oxidation and sulfate formation could be removed if there were a clear answer to this question since in a drop in equilibrium with a fixed SO2 concentration HSO; varies with [H’]-’ and SOi- with [H+]-‘. We will interpret the SO+xidation in terms of the more widely accepted view that the sulfite ion SO: - is the oxygen carrier and not HSO; or both. This assumption is in accordance for example with Winkelmann (1955) * krO is defined by: d[SO:-]/dt = klo [SO:-]. t In order to compare the results of Brimblecombe and their rate
pH 4 to 6
into our concept. $ For Penkett et 111.(1976) we calculated the corresponding k,, value for pH 6. 8 For Schroeter pH range 7-8
(1963) we calculated
k,.
values for the
i
Fig. 5. Pseudo first order rate constant for k,,,, as a function of pH.
According to Winkelmann (1955). the oxidatron follows the simple overall reaction.
PHASE
One of the most important questions in Sot-oxidation in the droplet phase is whether the bisulfite ion
Spedding (1974) with the others we transformed constants given in Table 1 of their paper for
In the following sections we wtll compare three dtfferent mechanisms of SO,-oxtdatlon and assess theu importance for sulfate formatron m the droplet phase In order to compare mvestrgattons. we have transformed all rate expressrons-whenever possible -Into the corresponding expresston of our concept wrth the sulfite ion (SO:-) being the oxygen earner.
reaction
10,
Most authors agree that the oxidation IS zero order with respect to oxygen over a wade range of 02-concentrations (for example Titoff. 1903; Riccoboni et al., 1949; Winkelmann, 1955; Penkett et al.. 1976). The order with respect to suUite is unity in most investigations (for example Titoff, 1903; Winkelmann, 1955; Rand and Gale, 1967). The following overall pseudofirst-order rate constants, k, 0, for reaction (10) are available* : (a) klo = 1.7 x 10-3s-’ at pH 6.8 and cn 25°C. Scott and Hobbs (1967). from measurements of Van den Heuvel and Mason (1963) at room temperature. between pH 2-4 at 25°C. (b) k,o = 3 x lo-‘s-r Miller and de Pena (1972). (c) klo = 3.5 x IO-‘s-l at pH 7 and 25°C. Winkelmann (1955). (d) k,,,t = 3.7 x iOT3-0.6 x 10-3s-’ between pH 4-6 at 25°C with 0.6 x lo-’ at pH 6. Brimblecombe and Spedding (1974). (e) k,J = 0.5 x 10e3 s-’ at pH 6 and 20°C. Pen kett ef al. (1976). (0 k,oij = C(I 6 x 10-3-0.6 x 10-3s-’ between pH 7-8 at 25°C. Schroeter (1963) (g) k,o = 1.2 x 10m4 [H+]-“.‘6 between pH 3-6 at 25°C. (k in s-l, [H’] in mol/l.). Beilke er al. (1975). (h) k,o = 0.013 + 59 [H+]‘!* (pH below 6 at 25°C) (k in s-‘, [H’] in mol/l.). MC Kay (1971) from measurements of Fuller and Crist (1941). Figure 5 shows klo values as a function of pH for 25°C. As seen in this figure, the k,,-values used by McKay (1971) and deduced from measurements of Fuller and Crist (1941) are higher than all other values by more than two orders of magnitude. One of the reasons for these high values is that Fuller and Crist did not take into consideration the variation of pH during the reaction nor the resulting changes of the ionic composition of the reaction mixture, even though such changes of the mole fraction of HSO; and SO:- are clearly indicated in Fig. 1
Heterogeneous
3
L
5
6
SO,-oxidation
PH
7
Fig. 6. Sulfate formation rates in the droplet phase as a function of droplet pH for three different SO,-oxidation mechanisms The calculated rates apply to the following conditions: gasphase concentrations for SOI: 1 ppb. O3 :40 ppb, T = 10°C.
For this reason we will confine ourselves to the other k,c-values given in the pH range 3 to 7. Although there is an opposite pH dependance of k,, between Beilke et al. (1975) and Brimblecombe and Spedding (1974). the values differ by no more than a factor of ca 2 in the pH range between 5 and 6. For pH 6 the corresponding value of .Penkett et al. (1976) is also within this range. Therefore it seems reasonable to accept a k10 value of ca 10-3s-’ in the pH range 5 to 6 at 25°C. Although the oxidation can be described by reaction (lo), the exact mechanism seems still to be a matter of discussion. It is likely that the oxidation proceeds via two chains with the radicals HO2 and SO; as chain carriers (Schmittkunz, 1963). Two results of Beilke er ol. (1975) who used a heterogeneous system to study SO,-oxidation should be mentioned here: Sulfate formation in aqueous solution was first order with respect to SO2 in the gas phase and did not depend on temperature within experimental error. This somewhat surprising result may well have arisen from the compensating effects of decreasing SO,-solubility (see Table I) and increasing reaction rate constant with increasing temperature. On the basis of this result an activation energy for reaction (10) has been estimated to be 14 kcal mol-’ in the pH range between 3 and 6 using the thermodynamic quantities given in Table 1. Since Beilke et al. (1975) used a heterogeneous system. their laboratory results will be directly applied to the atmosphere. For a cloud with a liquid water content (LWC) of 0.1 g rne3, S02-removal rates were calculated to fall between ca 10-6% h-i and 1.5% h-l for droplet pH values between 3 and 6, or ten times higher for LWC = 1 g rnb3. The lower dashed curves in Fig 6 show sulfate formation rates due to SO,-oxidation by O2 in the absence of catalysts as a function of droplet pH for 10°C and 1 ppb SO,-gas phase concentration. This SO,-gas phase concentration was
in the droplet phase
235
found to be+realistic at cloud level (1-5 km) in background air (Georgii and Jo& 1964; Gravenhorst. 1975). As seen in this figure, sulfate formation rates are extremely small except for high droplet pH (above 5 6) In Fig. 6 sulfate formation rates which were calculated on the basis of the experimental data of Brimblecombe and Spedding (1974) and of Penkett et al. (1976) are also included. A comparison of the three dashed curves shows that very similar conclusions concerning the role of SO,-oxidation by O2 in the absence of catalysts can be drawn. A special type of the SO,-oxidation by 0, in the absence of catalysts is the oxidation in the presence of ammonia (NH,). In the past, this mechanism has attracted substantial attention. It had previously been accepted by many scientists that the so-called SO,NHS-H,O-mechanism was the most effective SOroxidation process in atmospheric droplets. The e&t of ammonia is to keep the droplet pH high when S(IV) is oxidized to sulfate. This effect was investigated by many scientists, for example by Junge and Ryan (1958), Van den Heuvel and Mason (19633 Scott and Hobbs (1967). MC Kay (1971). Miller and de Pena (1972) Easter and Hobbs (1974) and Tomasi et al. (1975). The model of Scott and Hobbs was in best agreement with measurements by Beilke et al. (1975) It can, however, not be used (as for example Easter and Hobbs (1974) did) in order to calculate sulfate formation rates without further modifications of the physics involved (Beilke and Barrie, 1974) The calculated sulfate formation rates of Easter and Hobbs apply to a droplet pH of 6.6 (Harrison, 1974). It seems to be unlikely that such a high pH value is generally found. In a real atmosphere NH3 competes for control over droplet pH with many other soluble substances for example with an acidic background aerosol (Junge and Scheich, 1969; Gravenhorst. 1975) causing in most cases an overall droplet pH between ca 4-5. The sulfate formation rates in the droplet phase should therefore be lower than those predicted in the model of Scott and Hobbs by a factor between co 10’ to 10’ (Beilke, to be published). Therefore the major function of NH3 is likely to be no longer an effective enhancement of the SO,-oxidation but rather the conversion of a pre-existing sulfate containing droplet into a droplet containing ammonium and sulfate ions. However. the absorption of ammonia can shift the pH value from the 3 to 4 range to a range where oxidation of SO1 by O2 in the presence of catalysts is more important (45). If the source of sulfate in cloud drops is aqueous phase oxidation the oxidation mechanism must be effective between pH 3 and 5. 3.2 SO,-oxidation by 0,
in the presence oj catalysts
Catalysts are mainly transition metals of the 4th period especially manganese and iron. In the case of catalytic S02-oxidation by O2 in the presence of cata-
S. BEILKEand G.
736
lysts we will clearly separate aqueous phase SO,-OXIdation occurring in saline solutions of deliquesced aerosol salts with catalyst concentration of lo- ‘-1O-3 M from that occurrmg in dilute solutions like fog and cloud drops with catalyst concentrations between 1O-4-1O-8 M. We will confine ourselves to the latter. Since the important work of Bigelow m 1897 more than 200 publications have appeared in the literature dealing with catalytic oxidation of S(N). Most worked with S02- and catalyst concentrations which were higher by orders of magnitude than those occurring m the atmosphere. In addition, most studies were carried out with sodium sulfite solutions which have a pH range between 8-9. The atmospheric pH-range is between 3 and 6. Therefore we will restrict discussion to a few investigations carried out with concentrations of SO2 and catalysts similar to atmospheric conditions (catalyst concentrations between 1O-4-1O-s M and (or) SO,-concentration far below 1 ppm). The effect of catalysts was mvestigated using both a heterogeneous system (Barrie and Georgii, 1976) and a homogeneous one (Coughanowr and Krause, 1965; Bracewell and Gall, 1967; Tsunogai, 1971; Brimblecombe and Spedding. 1974; Penkett et al.. 1975; Lenelle. 1975; Betz. 1977). The oxidation of S(N) was found to be proportional to the square of the Mn2+-concentration when the concentration ratios S(IV)/Mn*’ 9 1. (Coughanowr and Krause, 1965) which is in agreement with the findings of Bracewell and Gall (1967). An empirical equation describing the oxidation in the presence of Mn” was given by Bracewell and Gall: B 1 A z = @iiy
+ [s(w)];.2
where A, B = constants, [Mn’+]e = mitial concentration of Mn SO., in aqueous solution, [S(IV)Je = Initial S(IV) concentration in aqueous solution and R = reaction rate. Btimblecombe and Spedding (1972) investigated the Fe3’-catalyzed SO,-oxidation in acidic aqueous solutions with concentration for Fe’+ = 10e6 M and S(IV) 2 lo-’ M. They found the disappearance rate of S(N) to be proportioned to both Fe3’ and S(IV). At a pH of 4 the oxidation proceeded faster in the presence of 10m6 M Fe 3 ’ than in the absence of catalysts by ca a factor 103. Barrie and Georgii (1976) investigated the catalyzed SO*-oxidation at 25°C and 8°C using single droplets of 2.1 mm diameter conaisting of distilled water with heavy metal concentrations in the range between low6 to IO-” molar for manganese and iron. The droplets were exposed to SO+oncentrations in air between 0.01-l ppm. The sulfate * For a droplet with 1Opm radius the calculated rate of SO,-diffusion to the droplet IS higher than the measured SOI-oxidation rate by a factor between ca IO’ to IO4 for the droplet pH range of interest here (3-5).
GRAVENHORST
formation rate was found to be first order with respect to SO2 in the gas phase. In the pH range 24 5 the most effective catalyst was Mn2’ followed hk Fe” and to a much lesser extent Fe’ + The catalytic effectiveness of Mn” was enhanced by the addition of Fe’+ or Fe3’ This synerglsttc effect was also observed by other authors (Penkett r~ 01. 1975. Lenelle, 1975). The Increase of temperature of a solution
of Mn’+ from 8 to 25°C caused an Increase in OXIdation rate by a factor 5 to IO m the pH range 2 4.5. However, the temperature dependance was reduced when Fe*’ or Fe3’-ions were added to the Mn’+ solution. On the basis of the experiments of Barrie and Georgia (1976), the conclusion can be drawn that sulfate formation due to catalytic SO,-oxidation can be significant in urban air clouds or fogs in which case the catalysts concentration can be of the order of 10m5 molar or even higher. A realistic approach with respect to concentrations and composition of catalysts is to use natural rainwater as solvent for studying the catalytic S02-oxidation. Betz (1977) has investigated the oxidation in a homogeneous aqueous phase of rainwater collected in the area of Frankfurt/M. (Germany). The heavy metal concentrations were between 10-7-l0-b molar for manganese and between 1O-6-1O-5 molar for iron. Of the total manganese SO-90% were &ssolved. For iron the corresponding fraction was 60-75%. The rainwater pH was between 3.2 and 5.2. By measuring the concentration decrease of S(N), sulfate formation rates were determined as a function of pH and temperature. The activation energy was measured to be 23 kcal mol- ’ which is in close agreement with the value of 21 kcal mol- ’ found by Penkett et cl. (1975). The rate constants are ca 1 to 2 orders of magnitude higher than those reported for the SO,-oxidation by O2 in the absence of metal catalysts. In order to apply the experimental data of Betz (1977) into the atmospheric heterogeneous system. three assumptions were made: (1) SOz-transport processes proceed fast compared with oxidation in the droplet phase.* (2) Sulfate formation proceeds first order with respect to SO2 in the gasphase. (3) The concentrations of heavy metals measured in rainwater are the same as in cloud or fog water. This assumption is to some extent uncertain. The results of our calculations (solid curve m Fig. 6) for a gas phase concentration of 1 ppb and 10°C show that sulfate formation proceeds faster than in the absence of metal catalysts by CO2 orders of magnitude. It has often been concluded on the basis of experiments with natural rainwater that this type of SO1-oxidation contributes only little to sulfate formation in the droplet phase at least for country clouds in rural areas. These conclusions are drawn under the assumption that catalyst concentrations and pH in rainwater are the same as in cloudwater which is highly uncertain. Very similar results have been
Heterogeneous
SO,-oxidation
obtained by Penkett et af. (1975a) and Lenelle (1975).* The oxidation rate measured by Tsunogai (1971) seems to be lower. The mechanism of the catalyzed %&oxidation by O2 is still a matter of discussion. The oxidation proceeds likely via two chains with the radicals HO* and SO, as chain carriers (Schmittkunz, 1963). The radicals are likely produced as follows: (see Barrie and Georgii (1976) who used the proposed m~h~isrn of ~hrni~k~ (1963)).
[MeX+ S(IV),J-‘6-“’ + 0,-r [Mer+ S(IV)~]-‘5-X’ + 0, 0;
+ H+-+HO*
[Me”’ S(IV)~]-‘5-X’ + SOi- - [Me”’ S(W),]-(“-*)
(14)
As can be seen, the complex [Me”’ s(IV),]-‘6-x) is regenerated Me ++ is transition metal ion of valency x. S(IV) is SO:- or HSO;. In the case of manganese as a catalyst, the results of Barrie and Georgii (1976) Seem to indicate that the manganese ion Mn 2+ formed complexes [Mn2+ (SO;-)s]-4 before taking part in the reaction. Interested readers are referred to the publication of Schmittkunz (1963) 3.3 SOL-ox~uriun by strongly oxidizing agem The effect of strongly oxidizing agents was investigated by Penkett (1972); Penkett and Garland (1974); Penkett, Jones and Eggleton (1975); Penkett, Brice and Eggleton (1976) and of Barrie (1975) for ozone (0,) and by Penkett, Brice and Eggleton (1976) for hydrogen peroxide (H202j The reaction of ozone with sodium sulfite in a homogen~us aqueous phase was studied by Penkett (1972) at pH 4.65 and was extended later to a pH range 1 to 5 by Penkett, Jones and Eggleton (1975). The SO,-oxidation was measured by following the decline in ozone. In these experiments the concentration of total sulfitet was kept in excess of ozone. In this case the reaction was first order with respect to ozone. The order with respect to total sulfite was unity at pH 4 and decreased with increasing PH. The O,-concentrations used were about 5 x tom5 M cor-
* Since the papers of Penkett er al. (1975) and of Lenelle (1975) are not yet published in the open literature, no further details on their work are gMn here. t Total sulfite is mainly HSO; in the range investigated by Penkett et al. @H 1-S). $ For a droplet with r = fO#rn radius the calculated rate of SOrdifFusion to the droplet is higher than the measured (Penkett et al., 1975) SO,-oxidation rate due to ozone by a factor between ca 10’ to 10’ for the droplet pH range 3-5. P Since the results of Pen&t, Brice and Eggleton (1976) have not been published yet in the open literature, no details are given here.
237
responding to a gas phase concentration of 4 2700 ppm in equilibrium. According to Penkett and Garland (1974) the results obtained with such high Osconcentrations can be extrapolated to atmospheric Os~onc~trations as indicated by their measurements on the rate of SO,-oxidation in a fog chamber using atmospheric concentrations of both ozone and SO*. The mechanism of the reaction between ozone and total sutfite is likely to proceed via a free radical mechanism (Penkett, Jones and Eggleton
(12) (13)
+ SO;
in the droplet phase
1975).*
It should be mentioned, however, that Barrie (1975) did not find a much enhanced SO+xidation rate due to ozone at a pH of co 4. He exposed three droplets (1.06mm radius) of distilled water to a stream of air containing a mixture of 0.4ppm SO, and different Os-concentrations between 0.01 and 0.5 ppm. In the droplet phase the increase of total sulfur was measured using a new isotope dilution technique (Klockow et al.. 1974). Barrie’s measurements with distilled water droplets showed only a very low enhancement of SO,-oxidation rate due to ozone. He found, however, an enhancement of the oxidation rate due to O5 when manganese was present. For an O,-gasphase concentration of 0.14 ppm the rate in a 10W5molar MnCl,droplet was doubled In order to estimate sulfate formation rates for the heterogeneous atmospheric droplet system on the basis of the experiments of Penkett (1972) and Penkett, Jones and Eggleton (1975) using a homogeneous aqueous phase. we used their kinetic data along with the thermodynamic quantities in Table 1 for S(W) and for ozone (Barrie, 1975j Again, we assumed transport processes for both SO1-gas and OS-gas to proceed fast compared with the SO,-oxidation1 in the droplet The uppermost dashed curve in Fig. 6 shows sulfate fo~tion rates due to SO,-oxidation by ozone after the experimental results of Penkett, Jones and Eggleton (1975) for the following set of conditions: gas phase concentrations are 1 ppb for SO, and 40 ppb for ozone and 10°C. The corresponding sulfate formation rate calculated on the basis of Barrie’s experiment is lower by a factor of co lo* at pH 4. Another oxidizing agent of importance for S02-oxidation in the droplet phase could be H20t for two reasons: (a) The solubility of H102 in water is extremely high. (b) The reported steady state concentrations of H20, in the lower troposphere are about 1 ppb (3 ppb according to model calculations of Levy (1973) and 1 ppb after Crutzen (1977) (persona1 communication to S. Beilke). Very recent experiments of Penkett, Brice and Eggleton (1976% have shown that the reaction is very fast even in the pH range 3 to 5. However, the importance of this oxidation process for sulfate formation in the droplet phase depends strongly on H202 gasphase con~~a~ons which have not yet been measured to our knowledge.
23X
5 BEILK~ and Ci. GRAVENHORST 4. IMPORTANCE SOI-OXIDATION
OF HFZTEROCENEOLS
IN THE ATMOSPHERE
FURTHER
AND
RESEARCH
Figure 6 shows a comparison of sulfate formation rates due to three S02-oxidation mechanisms (a) SO,-oxidation by O2 m the absence of catalysts (lower dashed curves). (b) S02-oxidation by O2 in the presence of catalysts (solid curve) and (c) S02-oxidation by ozone (upper dashed curve and one single dot) As seen in Fig. 6. oxidation type (a) IS unimportant for sulfate formation unless droplet pH is higher than ca 6. Although sulfate formation by type (b) proceeds faster by two orders of magnitude than bj type (a). it seems to be at least for background conditions of little importance for atmospheric sulfate formation unless droplet pH is higher than cu 5. For urban fog conditions, however. with higher heavy metal concentrations. this mechanism may play an important role. Sulfate formation rates due to S02-oxidation by ozone (type (c)) are higher than the corresponding rates for type (b) by ca 1 order of magnitude accordmg to Penkett er a/. (1975). In contrast. however, Barrle (1975) found that at least at pH 4 this rate is lower. The main discrepancies in heterogeneous SO,oxidation obvious from Fig. 6 are the pH dependance and the role of ozone. Therefore we suggest the following further research. (a) It should be determined whether the SOi--ion or the HSOJ-ion (or both) carries the oxygen in the pH range 3 to 7 This question is of fundamental Importance for the pH dependance of sulfate formation m atmospheric droplets. (b) Further research is necessary to explain the dlscrepancy of the experimental results of Penkett et al. (1975) and Barne (1975) concerning the role of ozone m heterogeneous SO,-oxidation in the absence and presence of catalysts. (c) In order to calculate sulfate formation rates due to SOI-oxidation by O2 in the presence of catalysts, we assumed that the concentrations of heavy metals and the pH value measured in rainwater are the same as m cloudwater. We therefore recommend measurements of the composltlon and pH of cloud and fog water since the small cloud and fog droplets are more important for sulfate formation due to their relatively long lifetime (minutes to hours) compared with the lifetime of large ramdrops (ca I minute). In a real atmosphere the three oxidation mechanisms proceed at the same time. We are therefore concerned with a very complex synergistic S02-oxidation by 02, 03. H,O, m the presence of NHJ, metal catalysts and organic compounds. Organic substances may either act as inhibitors or catalyst for SO,-oxidation and can increase or de* Droplet pH can easdy be measured with a micro-pH cell t Total sulfur m a smgle droplet can be measured with the Isotope dllutlon technique of Klockow YI al (1974)
crease droplet pH. Any ammonia absorbed by a droplet tends to increase the droplet pH which m turn causes higher sulfate formation rates due to oxidation by O2 in the presence of catalysts since the amount of S(IV) increases for a given SO,-gas phase concentration. In addition. the catalytic effectivity of any heavy metal of the 4th period depends also on pH. Furthermore. the catalytic effectivity of a special metal ion depends on the presence of other metals. The SO,-oxidation
rate obtained
with a mixture
of
iron IS. for example, higher than the sum of the rates with each ion present alone (Barne. 1975). In addition to the influence of pH the promoting effect of ozone for S02-oxidation, or the mhibiting effect of NO2 both m the presence of metal catalysts (Barrie. 1975) complicates the situation. (d) Because the overall effect of SO,-oxidation m droplets depends on the complex variety of interacbans, experiments should be carned out which take all these possibilities mto account. manganese and
These experiments should complement the studies suggested under (aHc) above. Such experiments should be carried out in a heterogeneous system. since m a homogeneous aqueous system compounds reacting in the liquid phase cannot be supplied from the gasphase to replace removed ones. That IS for example true for SOz. 03. H,02 and NH3. We suggest that drops consisting of natural rain- or cloudwater be suspended in outside air to simulate actual atmospheric conditions. By measuring the change of droplet pH* and increase of total sulfurt or sulfate, sulfate formation rates can be determined under realistic conditions. REFERENCES
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SO,-oxidation
Bracewell J. M. and Gall D. (1967) The catalytic oxidation of sulphur dioxide in solution at concentrations occurring in fog droplets Proc. Symp. on Physicochemical Transformation of Sulphur Dioxide Compounds in the Atmosphere and the Formation of Acid Smogs. Mainz. Germany, 1967. Brimblecombe P. and Spedding D. I. (1972) Rate of solution of gaseous sulphur dioxide of atmospheric concentrations. Nature 236, 225. Brimblecombe P. and Spedding D. 1. (1974) The catalyzed oxidation of micromolar aqueous sulphur dioxide-I. Atmospheric
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Crank I. (1976) The Mathematics of &&ion. 2nd edn. Clarendon Press, Oxford. Crutzen P. (1977) Personal communication to S. Beilke. Easter R. C. and Hobbs P. (1974) The formation of sulphates and the enhancement of cloud condensation nuclei in clouds. J. atm Sci. 31, 1586-1594. Eigen M., Kustin K. and Maass G. (1961) Die Geschwindigkeit der Hydratation von SOI in wPssriger LSsung. Z. Phys. Chem 30, 130. Fuller E. C. and Crist R. H. (1941) The rate of oxidation of sulfite ions by oxygen. J. Am. them. Sot. 63, lWl650.
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Klockow D., Denzinger H. and Riinicke R. (1974) Anwendung der substtihiometrischen Isotopenverdiinnungsanalyse auf die Bestimmung von atmosphlrischem Schwefel und Chlor in Background Luft Chem. Ing. Techn. 46. 18. Lenelle Y. (1975) Contribution a I’etude de l’oxydation des sulfites dans la pluie. Paper presented at the European sulfur symposium. Ispra, Italy, October 1975. Levy H. (1973) Photochemistry of minor constituents in the troposphere. Planet. Space Sci. 21, 575-591. Liss P. S. (1971) Exchange of SO2 between the atmosphere and natural waters Nature 233, 327-329. Liss P. S. and Slater P. G. (1974) Flux of gases across the air-sea interface. Natwe 247, 181-184. McKay H. A. C. (1971) The atmospheric oxidation of sul-
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Tsunogai S. (1971) Oxidation of sulfite in water and its bearing on the origin of sulphate in meteoric precipitation. Geochem. J. 5, 175. van den Heuvel A. P. and Mason B. J. (1963) The formation of ammonium sulfate in water droplets exposed to gaseous sulfur dioxide and ammonia. Q. JI. R. met. SOC. 89, 271-275. Wang J. C. and Himmelblau D. M. (1964) A kinetic study of sulfur dioxide in aqueous solution with radioactive tracers. AIChE Journal 21.402-404. Winkehnann D. (1955) Die elektrochemische Messung der Oxidationsgeschwindigkeit von Na2S0, durch geliisten Sauerstoff. Z. Eleklrochemie 59. 891-895.