SO2 oxidation by oxygen in monodisperse water droplets was studied in a cylindrical chamber, without and in the presence of ammonia. The range of SO2 concentration was from about 1022 to 5 ppmv, while the NH3 input concentration was kept constant at about 4.731022 ppmv. The contact time between gases and droplets was 210 s. The experimental results were compared with the theoretical

values predicted by the kinetics of Larson et al. (Atmos. Environ., 12 (1978) 1597) and McKay (Atmos. Environ., 5 (1971) 7). Much higher sulphate concentrations were

obtained in experiments run in the presence of NH3, as opposed to those without NH3. The experimental results agree with the values predicted by McKay’s kinetics and are higher than Larson’s

Prodi, F.

Santachiara, G.

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IL NUOVO CIMENTO Gennaio-Febbraio 1997VOL. 20 C, N. 1Oxidation of sulphur dioxide in water dropletsin the presence of ammoniaG. SANTACHIARA and F. PRODI (*)Istituto FISBAT-CNR - Bologna, Italy(ricevuto il 16 Settembre 1996; approvato il 5 Novembre 1996)Summary. — SO2 oxidation by oxygen in monodisperse water droplets was studiedin a cylindrical chamber, without and in the presence of ammonia. The range of SO2concentration was from about 1022 to 5 ppmv, while the NH3 input concentration waskept constant at about 4 .731022 ppmv. The contact time between gases anddroplets was 210 s. The experimental results were compared with the theoreticalvalues predicted by the kinetics of Larson et al. (Atmos. Environ., 12 (1978) 1597)and McKay (Atmos. Environ., 5 (1971) 7). Much higher sulphate concentrations wereobtained in experiments run in the presence of NH3 , as opposed to those withoutNH3 . The experimental results agree with the values predicted by McKay’s kineticsand are higher than Larson’s.PACS 92.60.Hp – Chemical composition and chemical interactions.1. – IntroductionSulphur dioxide, emitted from various sources such as fossil fuel combustion, thegas-phase of biogenic dimethylsulfides and vulcanic activity, can be oxidized in theatmosphere by homogeneous gas-phase reactions or heterogeneous processes, takingplace either on the surface of solid particles or inside water (cloud or fog) by means ofdissolved O2 , O3 or H2O2 . Trace substances frequently act as catalysts (Fe31 , Mn21 ,soot, etc.). The latter oxidation reactions are pH sensitive because the reaction ratesare influenced by [H1 ] or because the pH changes the solubility of SO2 .Sulphur dioxide oxidation can also be influenced by the presence of other gases,such as ammonia. For example, dry deposition of SO2 on vegetation is enhanced in thepresence of NH3 and with increasing air humidity [1].The role of NH3 in SO2 oxidation in an aqueous phase has been investigated in thepast by many scientists [2-5]. Junge and Ryan [6], in a pioneering paper on SO2oxidation in a liquid phase, measured a much larger amount of SO224 in the presence of(*) Also at: Dipartimento di Fisica dell’Università, Ferrara, Italy.111G. SANTACHIARA and F. PRODI112NH3 and were the first to observe the formation of ammonium sulphate in a muchshorter time than was expected from experiments in which bulk solutions of the sametype were mixed together, indicating that the rate of oxidation can be limited by themass-transport of gases.Easter and Hobbs [7] studied the SO2-NH3-liquid H2O system and calculated theproduction of ammonium sulfate by the oxidation of dissolved sulphur dioxide in clouddroplets in a wave cloud situation, following the Scott and Hobbs [3] oxidation rate. Theresults show that significant increases in the concentration of active CCN at 0.5%supersaturation can be produced in the oxidation of SO2 by the dissolved oxygen incloud droplets in a short time (4 min) and with concentrations of SO2 and NH3 typical ofunpolluted air (1 and 3 ppb, respectively).Field experiments also emphasize the importance of NH3 in the SO2 oxidationprocess. In a series of measurements in orographic clouds at Great Dun Fell(England), designed to investigate the oxidation of SO2 by means of H2O2 and O3 , Clarket al. [8] demonstrated the importance of the presence of NH3 , since the O3 reaction ispH dependent and the ammonia concentration has an important influence on the pH ofcloud water. The extrapolation of the results using a model demonstrates that, over awide range of SO2 concentrations, the total sulphate produced in the cloud is virtuallyindependent of the initial SO2 concentration and roughly proportional to theenvironmental ammonia concentration.It is important to assess the rate at which SO2 oxidation due to O2 may occur in theatmosphere, since the fastest rate observed in the laboratory [9] would make theprocess very important especially in the presence of NH3 , when compared withoxidation produced by other oxidisers in the atmosphere, while the lowest rate for thesame process [10] suggests that it is of little consequence.To increase our knowledge of the influence of ammonia in the oxidation process ofSO2 , we carried out laboratory experiments, using monodispersed droplets. These wereused instead of polydispersed droplets, since the gas and liquid mass-transfercoefficients depend on the radius of the droplets. Using the latter, the SO2 oxidationrate could have depended on the droplet size distribution.2. – ExperimentalThe experimental techniques used in these experiments have already beenpresented in detail in [11]. A stream of air is passed through a Collison generator,which atomizes a dilute NaCl aqueous solution, thus producing condensation nuclei.The gas flow at the outlet is partly sent to a thermostated bath, where the water isvaporized at T480 7C. The droplets formed on condensation nuclei are vaporized and acontrolled recondensation produces a cloud of monodispersed droplets. The dropletcharacterization technique employed a glass slide coated with black carbon andthermophoretic enhancement of droplet deposition by cooling the slide. The ratio of thedroplet to spot diameter is 0.86. The droplet radius was R4 (2 .060.3) mm.At the generator outlet the droplets enter a reaction chamber togheter with amixture of SO2 and NH3 . Air, SO2 and NH3 are supplied by compressed-gas cylindersand their flow-rate ratio is varied to obtain variable concentration. The total flow was(32565) l h21 . The chamber employed consists of a pipe measuring 200 cm in lengthand 11 cm in diameter and was heated slightly on the outside to prevent condensationon the inside surface. The contact time is 210 s. At the outlet of the pipe, the dropletsOXIDATION OF SULPHUR DIOXIDE IN WATER DROPLETS ETC. 113are captured on a specifically designed, circular trough impactor with a near unitarycapture efficiency. The experiment lasts about 30 min. The gas stream at the outlet ofthe impactor is sucked for about 158 through a bubbler containing 40 cc of a 5% H2 O2and afterwards, for the same time period, through 40 cc of 0.1 N H2 SO4 to measure theSO2 and NH3 concentrations in the gas.The chemical analyses of SO224 , NH14 and Cl2 are performed with a Dionex ionchromatograph. Liquid water content (LWC), determined by the ratio between thewater mass collected during each run and the flowed gas volume was (3 .860.1) g m23.3. – Theoretical and experimental resultsThe experiments are made using monodispersed droplets grown on NaCl condens-ation nuclei at T4 (2561) 7C. Cl2 concentration in the droplets is (4.3 6 2.2) ppm.The experimental data are reported in figs. 1 and 2. For comparison we report dataobtained from the kinetics proposed by McKay [12], deduced from measurements ofFuller and Crist [9], and Larson et al. [13], for SO2 oxidation due to oxygen in thepresence of ammonia and in the absence of a catalyst. These researchers suggest thatFig. 1. – Sulphate concentration in the liquid phase as a function of SO2 concentration in the gasphase, in the absence of NH3 . Solid line, McKay’s oxidation rate; dashed line, Larson’s et al.oxidation rate. Squares represent average of three to five runs and bars indicate standarddeviation.G. SANTACHIARA and F. PRODI114Fig. 2. – Sulphate concentration in the liquid phase as a function of SO2 concentration in the gasphase, in the presence of NH3 . Solid line, McKay’s oxidation rate; dashed line, Larson’s et al.oxidation rate. Squares represent average of three to five runs and bars indicate standarddeviation.the rate of production of SO224 is given by the first-order rate equationd[ SO224 ]Odt4k[ SO223 ] .(1)The rate constants proposed by McKay [12] and Larson et al. [13] at T4298 K are,respectively,k140.013159 [H1]0.5 s21 , k244.8 Q102318.9 [ H1 ]0.513.9 Q10212 p02 [ H1 ]21 s21 ,where p0240.2 atm is the partial pressure of oxygen in the gas, [H1 ] is theconcentration of hydrogen ion in the liquid phase, M. We choose these rate constantsfor comparison with the experimental data, as McKay’s kinetics give a higherconversion rate with respect to all the values proposed [3, 10, 14, 15], while Larson etal. [13] give lower values than McKay’s oxidation rate by almost one order ofmagnitude.Chemical equilibrium equations, equilibrium constant expressions and values aregiven in table I.By considering also the electroneutrality equation,[ H1 ]1 [ NH14 ]4 [ OH2 ]1 [ HSO32 ]12[ SO223 ]12[ SO224 ] ,OXIDATION OF SULPHUR DIOXIDE IN WATER DROPLETS ETC. 115TABLE I. – Reactions and equilibrium constants (from Seinfeld [16]).Chemical equilibriumequationEquilibrium constantexpressionValue of the equilibriumconstant at T4298 KH2 O4H11OH2 Kw4 [ H1 ] [OH2 ] 1 .008310214 M2( SO2 )gas1H2 O4SO2 QH2 O Khs4 [ SO2 QH2 O ]OpSO2 1.24 M atm21SO2 QH2 O4H11HSO23 Ks14 [ H1 ]1 [ HSO32 ]O[ SO2 QH2 O ] 1 .2931022 MHSO324H11SO43 Ks24 [ H1 ] [SO34 ]O[ HSO32 ] 6 .01431028 M( NH3 )gas1H2 O4NH3 QH2 O Kha4 [ NH3 QH2 O ] 62 M atm21NH3 QH2 O4NH14 1OH2 Ka14 [ NH41 ] [OH2 ]O[ NH3 QH2 O ] 1.70931025 Mwe obtain the following equation:H3 (11Kha Ka1 OKw ) pNH322H2 [ SO224 ]2H(Kw1Khs Ks1 pSO2 )22Ksh Ks1 Ks2 pSO240 .In our calculations (at T4298K ) we suppose a constant LWC and an instantaneousequilibrium between the gas and the liquid phase for each gas considered( O2 , SO2 , NH3 ). Therefore diffusion is not considered to be a controlling factor, eitherin the gas or the liquid phase and the only chemical step is the sulphate oxidation. Frominput SO2 , NH3 gas-phase concentration and LWC, an iterative program calculates theaqueous and the residual gas-phase concentrations that result when the SO2 and NH3are dissolved in liquid water. The equilibrium between the two phases produces adepletion of NH3 and SO2 in the gas phase. While for input SO2D0.5 ppmv depletion inthe gas phase is small and decreases with increasing SO2 concentration, for NH3 it isalways appreciable with respect to the initial concentration. For example, with LWC43.8 g m23, NH340.047 ppmv and SO240.5 ppmv, the equilibrium value for NH3 andSO2 are, respectively, 0.34 ppbv and 0.452 ppmv. So, at equilibrium, we have ammoniagas-concentrations comparable with those present in the atmosphere.We do not consider the CO2 equilibrium, since in our experiment the water used toproduce monodisperse droplets is at t`80 7C, so it contains no dissolved CO2 .The initial rate of production of SO224 is given by (1). Sulphate production changesH1, NH14 , and SO223 concentration in the liquid phase and NH3 and SO2 equilibriumconcentrations in the gas phase.Therefore, we can calculate sulphate production with an iterative procedure afterthe contact time t4210 s used in our experiments, to compare the theoretical data withthose obtained experimentally. The input NH3 concentration (before transfer of NH3from a liquid to a gas phase) is (0.04760.03) ppmv. Input NH3 (as for SO2 )concentrations were measured from NH14 (SO224 ) measured in the captured dropletsand from NH3 (SO2 ) measured in the gaseous phase, at the outlet of the column.In fig. 1 we also report sulphate production due to the oxidation of O2 , in theabsence of ammonia, following the same kinetics, and experimental results.G. SANTACHIARA and F. PRODI1164. – DiscussionFrom fig. 1 we can observe that sulphate concentration in a liquid phase, producedafter 210 s, following McKay and Larson’s kinetics, increases slowly with increasing SO2gas concentration in the absence of ammonia. This depends on the fact that Larson andMcKay’s kinetics suppose that the SO223 ion is oxidized and the theoretical equilibriumshows that SO223 concentration changes are small in relation to the SO2 concentration.For SO2 concentrations comparable with the atmospheric values, the calculationsshow a sharp increase in sulphate production in the presence of ammonia (NH340.047 ppm) , compared to when it is absent. Sulphate concentration shows a maximumwhen SO2 and NH3 input concentrations are comparable (fig. 2). In fact, ammonia canbuffer the solution pH, so there is an increase of SO223 . A further increase of SO2concentration can reduce sulphate production due to the limited buffering capacity ofthe ammonia. That is, while an increase of ammonia concentration should theoreticallyalways produce an increase in sulphate production, this is not the case for SO2 .By examining our laboratory results, we observe that experimental data can beexplained by supposing that the SO223 ion is oxidized, following McKay’s kinetics and inabsence of diffusive resistence. In fact, following Freiberg and Schwartz [17] andconsidering the droplet radius and gas concentrations used in our experiments, thecharacteristic time for the gas (t d, g ) and the liquid-phase diffusion (t d, a ) are muchlower than the characteristic time associated with the oxidation reaction in the liquidphase, relative to the gas-phase SO2 concentration (t c , g ) or to aqueous sulphur-IVspecies (t c , a ), respectively. For example t d, g is about 1028 s for all gases considered( O2 , SO2 , NH3 ) and t d, a`1024 s in liquid phase. If we consider for example an inputSO2 and NH3 gas-phase concentration of 231022 and 4 .731022 ppmv, respectively,following McKay’s kinetics, the t c , a will range from 60 to 120 s (during a 210 s period oftime) and t c , g will range from 231025 to 1024 s. In addition, with the same SO2 andNH3 concentrations, the characteristic time (treag ) required for gas-phase diffusion tosupply an amount of reagent (SO2 ) to the droplet necessary to produce a concentrationof dissolved reagent in equilibrium with its concentration in the bulk gas-phase, isabout 0.5 s (in the absence of NH3 it would be about 1023 s) which is much lower thanthe experimental duration (210 s). Characteristic times of the aqueous dissociationreactions of SO2 are very short when compared with the others. Consequently, we mustconsider the several aqueous sulphur-IV species to be in equilibrium.A discrepancy exists in the available literature on the aqueous oxidation of SO2 byoxygen. Many researchers [3, 15, 18] agree that the uncatalyzed oxidation of SO2 inaqueous systems occurs by a first-order reaction in SO223 and no oxidation of HSO23occurs directly, but the rate constants proposed are very different. For example, atpH46 and T4298 K, Fuller and Crist give kB731021 s21 , while Brimblecombe andSpedding 631023 s21 . Thus disparity may be due to impurity in either the reagent orthe water or to mass-transport limitation (in the gas and/or aqueous phase). FollowingDasgupta [19], probably the rate observations most easily reproduced were made byFuller and Crist.Some researchers propose further kinetics. Hansen et al. [20] studied the SO2oxidation in the presence of NH3 in laboratory clouds, by using NaCl and soot particlesas cloud condensation. The rate of SO224 formation was observed to be as much as twoorders of magnitude greater than that predicted by Larson’s equation. According tothese researchers sulphate formation is essentially independent of the type of nuclei onwhich the cloud forms.OXIDATION OF SULPHUR DIOXIDE IN WATER DROPLETS ETC. 117Benner et al. [21], in a laboratory experiment in a thin water film, found that theoxidation rate exceeded the rate expression of Larson et al. by about one order ofmagnitude, i.e. a smaller enhancement with respect to that observed in the previouslyreported droplet reaction. Hansen et al. and Benner et al. interpreted their findings interms of the enhanced adsorption of the reacting gases on the droplet surface, asreported by Jayne et al. [22]. These authors suggest that at the gas-liquid interface,gaseous SO2 collides with a water molecule and forms the surface complex HSO23 -H1(chemisorption process) which is in equilibrium with the gas-phase SO2 . Such anintermediate surface complex may then enter the bulk-liquid as HSO23 .Adema and Heeres [23] in a study on the dry deposition of SO2 in air on a thin waterlayer (4–4600 mm), in the presence of ammonia , concluded that the oxidation of HSO23occurs at the surface of the water and that the results cannot be explained by theoxidation of the SO223 ion.As regards Hansen’s results, there is important evidence that cloud formationoccurs in the presence of reactive gases, so initially there were much highconcentrations of catalyzers (Cl2 and soot) in the droplets, that can influence theoxidation process. This condition could at least partially explain the high sulphateconversion obtained. Moreover, Jayne et al. [22], by studying the effects of the H2 O2content on the uptake of SO2 by aqueous droplets, concluded that there is nomeasurable reaction of H2 O2 with surface chemisorbed SO2 , i.e. the bulk reactionkinetics of H2 O2 with dissolved S(IV) dominate the reactive uptake. Thus, evensupposing the formation of a surface-complex, this would not necessary produce anincrease in the oxidation of SO2 in liquid droplets. In conclusion, according to ourexperimental data, the oxidation of SO2 in cloud droplets, with or without ammonia,could be explained by McKay’s kinetics, with no need to introduce the assumptionsmade by the researchers quoted above.* * *The cooperation of Area Analitica Acque - Settore Chimico - P.M.P. - AziendaU.S.L., Bologna, Italy in performing the chemical analyses, D. SANTACHIARA incarrying out the computations and the assistance of M. TERCON are gratefullyacknowledged.R E F E R E N C E S[1] VAN HOVE L. W. A., ADEMA E. H., VREDENBERG W. J. and PIETERS G. A., Atmos. Environ.,23 (1989) 1479.[2] VAN DEN HEUVEL A. P. and MASON B. J., Q. J. R. Meteorol. Soc., 89 (1963) 271.[3] SCOTT W. D. and HOBBS P. V., J. Atmos. Sci., 24 (1967) 54.[4] MILLER J. M. and DE PENA R. G., J. Geophys. Res., 77 (1972) 5905.[5] TOMASI C., GUZZI R. and VITTORI O., J. Atmos. Sci., 32 (1975) 1580.[6] JUNGE E. and RYAN T. G., Q. J. R. Meterol. Soc., 84 (1958) 46.[7] EASTER R. C. and HOBBS P. V., J. Atmos. Sci., 31 (1974) 1586.[8] CLARK P. A., GERVAT G. P., HILL T. A. and MARSH A. R. W., J. Geophys. Res., 95 (1990)13985.[9] FULLER E. C. and CRIST R. H., J. Am. Chem. Soc., 63 (1941) 1644.[10] BRIMBLECOMBE P. and SPEDDING D. J., Atmos. Environ., 6 (1974) 937.[11] SANTACHIARA G., PRODI F. and VIVARELLI F., Atmos. Environ., 23 (1989) 1775.G. SANTACHIARA and F. PRODI118[12] MCKAY H. A. C., Atmos. Environ., 5 (1971) 7.[13] LARSON T. V., HORIKE N. R. and HARRISON H., Atmos. Environ., 12 (1978) 1597.[14] SCHROETER L. C., J. Pharm. Sci., 52 (1963) 559.[15] BEILKE S., LAMB D. and MULLER J., Atmos. Environ., 9 (1975) 1083.[16] SEINFELD J. H., Atmospheric Chemistry and Physics of Air Pollution (Wiley, New York)1986.[17] FREIBERG J. E. and SCHWARTZ S. E., Atmos. Environ., 7 (1981) 1145.[18] RADOJEVIC M., Environ. Technol. Lett., 5 (1984) 549.[19] DASGUPTA P. K., Atmos. Environ., 14 (1980) 272.[20] HANSEN A. D. A., BENNER W. H. and NOVAKOV T., Atmos. Environ. A, 25 (1991) 2521.[21] BENNER W. H., OGOREVC B. and NOVAKOV T., Atmos. Environ. A, 26 (1992) 1713.[22] JAYNE J. T., DAVIDOVITS P., WORSNOP D. R., ZAHNISER M. S. and KOLB C. E., J. Phys. Chem.,94 (1990) 6041.[23] ADEMA E. H. and HEERES P., Atmos. Environ., 29 (1995) 1091.

Oxidation of sulphur dioxide in water droplets in the presence of ammonia

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Oxidation of sulphur dioxide in water droplets in the presence of ammonia

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