Modeling Sulfur Dioxide Deposition on Calcium Carbonate - American

Feb 7, 2003 - Juan Luis Pe´rez Bernal and Miguel Angel Bello*. Departamento de ... nesses and then drop from the stone surface.6 The stone surface wh...
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MATERIALS AND INTERFACES Modeling Sulfur Dioxide Deposition on Calcium Carbonate Juan Luis Pe´ rez Bernal and Miguel Angel Bello* Departamento de Quı´mica Analı´tica, Facultad de Quı´mica, Universidad de Sevilla, 41080 Sevilla, Spain

Sulfur dioxide is the main decay factor of carbonate-based stones (limestones, marble, etc.) in polluted environments, with crust formation and solubilization of the stone being the main components. The deposition and presence of airborne particles can accelerate and enhance the decay processes due to sulfur dioxide. Different laboratory tests have been carried out to study the deposition and reaction of sulfur dioxide with calcium carbonate. The rate and extension of reactions is strongly influenced by the relative humidity and the presence of foreign substances. It has been found that the evolution of the concentration of the involved species (CaCO3, CaSO3‚ 1/ H O, and CaSO ‚2H O) can be represented by a mathematical model of the form X ) k [1 2 2 4 2 1 exp(k2t)], where X is the concentration expressed as a molar fraction, t is time, k1 is the concentration as t tends to infinity, and k2 can be associated with a pseudo-first-order rate constant. The effect of different metallic oxides and salts, silica gel, and activated carbon has been tested. 1. Introduction The adverse effect of sulfur dioxide on limestones, which are widely used in historical buildings, is wellknown. Sulfur dioxide, in addition to atmospheric particulate matter and water (liquid or gaseous), is accepted as the main decay factor in urban and polluted environments.1-3 Almost all of the studies carried out tend toward the establishment of the origin of gypsum detected on the surface of weathered materials, elucidate the formation mechanism, and identify weathering factors involved in the decay processes. The main effects of sulfur dioxide on limestones are the formation of crusts and the loss of material due to solubilization, which can represent the 30-50% of material lost.4,5 The loss of material can also be produced where weathering crusts reach certain thicknesses and then drop from the stone surface.6 The stone surface where the crust is detached usually presents disaggregation and higher porosity and surface area than the original stone, being then weaker to further weathering processes.7 If previous formation of calcium sulfite was considered, the formation of gypsum from the reaction of calcium carbonate with sulfur dioxide can be schematized as SO2

CaCO3 9 8 CaSO3‚1/2H2O H O 2

O2

CaSO3‚1/2H2O 9 8 CaSO4‚2H2O H O 2

Another possibility is the absorption of sulfur dioxide * Corresponding author. Tel.: +34 (9)54557172. Fax: +34 (9)54557168. E-mail: [email protected].

in rainwater, liquid atmospheric aerosols, or moist film supported on a stone surface, where it is oxidized to form a sulfuric acid solution that dissolves the calcium carbonate by gypsum formation: O2

SO2 9 8 H2SO4 H O 2

H 2O

CaCO3 + H2SO4 98 CaSO4‚2H2O The presence of at least a minimum water content is critical so that the processes involved occur.8-11 The water, liquid, or gas has a great influence on the material or surface reactivity and on the reaction rate, with those ones with the water content growing.9,10,12 It has been observed that there is no reaction with sulfur dioxide in an environment with a relative humidity (RH) lower than 40%.8,9,12 Water can be present physically absorbed, condensed in micropores, and as liquid water in the material voids; in the last situation, it has been observed that the sulfur dioxide deposition is greatly favored.12,13 Some authors consider that only the presence of a minimum quantity of water is necessary for the oxidation of sulfur dioxide and gypsum formation;10,14,15 on the other hand, others consider that the presence of a catalyst is also necessary so that the reactions can proceed.8,16-19 The number of works devoted to the study of the mechanism of sulfur dioxide oxidation and gypsum formation in relation to stone decay and the influence of stone foreign matter is rather poor. The majority of works are devoted to the comparison of the reactivity of different stones or to the establishment of the effectivity of protective treatments. In addition, the results are very dependent on the experimental setup, so it is difficult to compare them.

10.1021/ie020426h CCC: $25.00 © 2003 American Chemical Society Published on Web 02/07/2003

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Figure 1. Climatic chamber block diagram.

Elfving8 has studied the catalytic effect of iron(III) and mangenese(IV) oxides on the sulfur dioxide and calcium sulfite oxidation using pure calcite as the reaction substrate, with mangenese(IV) oxide being a better catalyst than iron(III) oxide and the quantity of gypsum formed being dependent on the catalyst concentration. Hutchinson et al.20 have reported the catalytic activity of copper(II), iron(III), and mangenese(IV) oxides and fly ash in the oxidation of sulfur dioxide on pure calcium carbonate and limestone samples; while in limestone samples, it seems that the catalyst presence has no effect, in the calcium carbonate samples, its presence is very important. With respect to fly ashes, they have observed that in the presence of calcium chloride they have a clear catalytic effect. Ausset et al.21,22 have also studied the effect of fly ashes, and they have observed that they favor the formation of gypsum, that they appear to be intimately bonded to fly ash particles, and that the deposition of sulfur dioxide tends to decay with time.22 Sabbioni et al.23 have studied the influence of carbonaceous particles on the reactivity of carbonatic stones, and they have observed that particles increase the gypsum formation, especially those with a high heavy-metal content. Bo¨ke et al.24 have studied the effect of different substances, montmorillonite, iron(III) and vanadium(V) oxides, activated carbon, and copper(II) chloride, on the reactivity of calcium carbonate. They reported that montmorillonite seems to have no effect on the reactivity of calcium carbonate and that cupric chloride and metallic oxides enhance the reactivity of calcium carbonate. Up to now, there is no proposal of a mathematical model to characterize the evolution of gypsum and/or calcium sulfite formation. In this work a mathematical model is proposed, and it allows one to characterize the effect of foreign substances on the reactivity of calcium carbonate and to make comparisons between possible catalysts. It also makes it possible to evaluate reaction rates as well as to establish the effect of the concentration of foreign substances. Future work will establish and compare thoroughly the effect of the different substances studied as possible catalysts. 2. Methodology 2.1. Climatic Chamber. A homemade climatic chamber25-28 was used for all of the tests. Some modifications on the first design were carried out to improve its versatility and usefulness. Figure 1 shows the schematic diagram of a chamber design. An air compressor is used to carry the gas

supply. The air current is forced to pass through water and particle traps before entering the drying and purifying sections. The drying section consists of two parallel lines equipped with concentrated sulfuric acid drying towers and aerosol traps. The purifying section consists of two parallel lines equipped with a silica gel with an indicator trap that is used as the indicator of the drying section, a soda lime trap for the retention of acid aerosols, and an activated carbon trap for the retention of organic aerosols and vapors. After the drying and purifying sections, the air current is divided in two separate lines equipped with flowmeters and regulators: the first one is the dry air current, and the second one is passed through thermostatized deionized water. Both air currents are mixed in a mixing chamber and passed through a thermohygrometer before entering the exposition chamber. All tests have been performed at constant RHs (60% and 80%), avoiding condensation. The pollutant (sulfur dioxide) is added from a gas cylinder (supplied by Air Liquide) with a purity higher than 99.9%, equipped with a manometer, a micrometric needle, and a capillary flowmeter. Both the sulfur dioxide and the air current are mixed in a thermoestatized mixing chamber. All air and pollutant conductions are thermoestatized to the working temperature (40 °C). Air flow in outdoor environments is greater than that in laboratory-exposure systems. Accordingly, to maintain the same pollutant presentation rate, a higher concentration of pollutant should be used.29 Considering an average sulfur dioxide concentration of 60 µg m-3 and a wind speed of 2 m s-1 results in a presentation rate of the pollutant of 120 (µg of SO2) m-2 s-1. All tests have been performed using a sulfur dioxide concentration of 1000 ppm. Taking into account the sample area (19 cm2) and the air flow (0.1 L s-1), this leads to a presentation rate of the pollutant of 130 (µg of SO2) m-2 s-1, which is similar to the average presentation rate for outdoor environments.29 The sulfur dioxide concentration is monitored using a continuous sulfur dioxide analyzer (The Analytical Development Co., ADC, model RF2B), with a working range from 0 to 2000 ppm. For the calibration of the analyzer, a gas cylinder of 1015 ((0.1) ppm of sulfur dioxide in air is used (supplied by Air Liquide). Variation of the sulfur dioxide concentration in the test chamber is lower than 5% of working conditions. 2.2. Sample Preparation. All reagents used as a possible catalyst and sample matrix were of analytical quality. Powdered calcium carbonate (Merck), in the form of precipitated calcite, was used as a sample matrix. Substances tested as catalysts were metallic oxides, iron(III), mangenese(IV), nickel(II), copper(II), lead(II), zinc(II), vanadium(V), and calcium(II) chlorides, activated carbon, and silica gel (Merck). Possible catalysts were added to calcium carbonate at a concentration of 1% (w/w). For calcium(II), copper(II), mangenese(II), nickel(II), iron(III), and zinc(II) chlorides, mangenese(IV) oxide and activated carbon samples with 5%, 10%, and 20% (w/w) were also tested. Samples were mixed with water, in a proportion of 1:1 (w/w), to form a homogeneous suspension that was dispersed between two microscopic slides (7.6 cm × 2.6 cm) and allowed to dry for 48 h at 40 °C. This procedure allows one to obtain very thin and reproducible samples, with a weight-to-area ratio between 0.006 and 0.008 g cm-2.

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KBr pellets, using a Perkin-Elmer Paragon 500 FTIR. Figure 2 shows the infrared spectra for the compounds determined; the peaks used for the quantitation correspond to the C-O bend tension at 1453 cm-1 for calcium carbonate and the S-O bend tension at 1146 and 980 cm-1 for calcium sulfate and sulfite, respectively. All reagents used for the calibration were of analytical quality (Merck). As calibration standards, different mixtures of calcium carbonate, calcium sulfate, and calcium sulfite have been prepared. Calcium sulfite was sintetized from a 0.1 M calcium chloride dissolution and a 0.1 M sodium sulfite dissolution; the purity of the obtained calcium sulfite was tested using infrared spectroscopy and X-ray diffraction. For calibration a modified internal standard technique was followed. The peak of calcium carbonate was taken as an internal reference; thus, the ratio of calcium sulfate or sulfite absorbance with respect to calcium carbonate absorbance was represented vs the mole ratio of calcium sulfate or sulfite with respect to calcium carbonate (RCaSO4 and RCaSO3), Figure 3 shows the calibration graph and the residual plots obtained by this method. Once the mole ratios are known, the mole fractions (XCaCO3, XCaSO4, and XCaSO3) can be easily calculated from

XCaCO3 ) Figure 2. Calcium carbonate, calcium sulfate dihydrate, and calcium sulfite hemihydrate infrared spectra.

2.3. Analytical Methodology. The determinations of calcium carbonate (C), calcium sulfate dihydrate (G), and calcium sulfite hemihydrate (S) have been carried out using Fourier transform infrared spectroscopy in

XCaSO4 ) XCaSO3 )

1 1 + RCaSO4 + RCaSO3 RCaSO4 1 + RCaSO4 + RCaSO3 RCaSO3 1 + RCaSO4 + RCaSO3

Figure 3. Calibration and residual plots for calcium sulfate dihydrate and calcium sulfite hemihydrate.

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Figure 4. Evolution of samples with CaCl2 exposed to SO2 (1000 ppm) in air at 40 °C and RHs 60% and 80%.

3. Results and Discussion Two different tests have been performed. In the first run, samples have been exposed to 1000 ppm of sulfur dioxide in air at 40 °C and 60% RH up to a maximum time of 1500 h. Second run series have been performed with the same pollutant concentration and a RH of 80% and with a maximum exposure time of 240 h. The sample reactivity was greatly influenced by the RH. In almost all cases, at the end of exposure time, the concentrations of products were higher in the samples exposed to 80% RH. In the few cases where the final composition of samples did not differ greatly between them, a faster reaction has been observed for the samples exposed to 80% RH. The tested samples showed different behaviors depending on the substance added. Pure calcium carbonate samples (blank samples) exposed to 60% RH did not react after 1500 h of exposure. Blank samples exposed to 80% RH show the formation of a small amount of calcium sulfite, lower than 0.1 mole fraction, and it was possible to detect it after 12 h of exposure time. Figure 4 shows the evolution of the composition of samples with CaCl2 (1% w/w) exposed to 60% and 80% RH. The calcium sulfite concentration is greater at the end of the experience carried out with 80% RH. In addition, calcium sulfite formation is faster for the samples exposed to 80% RH. It is also remarkable that in the second sample set it is possible to detect the formation of small amounts of gypsum, while in the samples exposed to 60% RH, there is no evidence of gypsum formation. Figure 5 shows the evolution of samples with MnCl2 (1% w/w) exposed to 60% and 80% RH; as can

Figure 5. Evolution of samples with MnCl2 exposed to SO2 (1000 ppm) in air at 40 °C and RHs 60% and 80%.

be seen, the samples are, again, much more reactive in the run carried out with 80% RH and, in addition, MnCl2 shows a clear catalytic activity with respect to S(IV) to S(VI) oxidation; in both cases calcium sulfite has not been detected. Figure 6 shows the evolution of samples with CuCl2 (1% w/w) exposed to 60% and 80% RH. In this case it is possible to appreciate the formation of calcium sulfite as well as calcium sulfate. All of the samples showed one of the next behaviors: formation of calcium sulfite, formation of calcium sulfate, and formation of both products. In addition, the evolution of composition with respect to time also showed the same tendency in almost all of the samples: first a fast growth of product concentration and then a slower tendency to a maximum value. The reaction steps are supposed to be30-32 (1) diffusion of sulfur dioxide to the sample surface, (2) adsorption of sulfur dioxide, (3) oxidation S(IV) f S(VI) (depending on conditions), and (4) reaction with calcium carbonate to form calcium sulfate and/or calcium sulfite. Chun and Quon30 proposed the capacity-limited heterogeneous reaction model. That model assumes that active sites of the sample become occupied by the reaction products and are not available for further reaction. So, the reaction of sulfur dioxide over those particles is expected to reach a maximum depending on the number of active sites available. Also, the covering of the original calcium carbonate by the reaction products and the passivation of the calcium carbonate surface it must be considered for subsequent reaction.33,34 The rate-controlling step is supposed to be the adsorption of sulfur dioxide over the sample surface

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Figure 7. Calcium sulfite concentration (mole fraction) evolution for samples with CaCl2 (1% w/w) exposed to 80% RH, model CaSO3 3 adjustment (dashed line), and residual plot (δ ) Xmodel - XCaSO exp ). Figure 6. Evolution of samples with CuCl2 exposed to SO2 (1000 ppm) in air at 40 °C and RHs 60% and 80%.

(step 2). If Θ0 is the fraction of free active sites and Θ is the fraction of occupied sites at time t, assuming firstorder kinetics, the adsorption rate of sulfur dioxide can be expressed by

dΘ/dt ) kPΘ0 where P is the sulfur dioxide concentration, which was held constant, t is the time expressed in hours, and k1 is a rate constant. Substituting Θ0 by 1 - Θ and integrating, we arrive at

Θ ) 1 - e-kPt Expressing kP as k2 and multiplying by the maximum capacity of sulfur dioxide adsorption (k1), we arrive at

Xi ) k1(1 - e-k2t) where Xi is the calcium sulfate or sulfite mole fraction formed at time t. The time evolution of the composition of exposed samples, expressed as mole fraction, follows the above expression, where Xi is the mole fraction of the species considered (calcium sulfite or calcium sulfate), k1 is an adjustable parameter and equals the mole fraction of the product considered as time tends to infinity, k2 is the second adjustable parameter and can be associated with a pseudo-first-order velocity constant, and t is time expressed in hours. Some relevant results are described concisely below as examples of model application.

Figure 8. Calcium sulfate concentration (mole fraction) evolution for samples with MnCl2 (1% w/w) exposed to 60% RH, model CaSO4 4 adjustment (dashed line), and residual plot (δ ) Xmodel - XCaSO exp ).

Figure 7 shows the evolution of the calcium sulfite concentration (in mole fraction) of samples with CaCl2 (1% w/w) exposed to 80% RH and the adjusted model,

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Figure 9. Sample composition evolution for samples with FeCl3 (1%, 5%, 10%, and 20% w/w) exposed to 80% RH and model adjustment (dashed line).

as well as the residual plots for model adjustment. The resulting parameters are k1 ) 0.40 and k2 ) 0.0082, with a correlation coefficient r ) 0.9918. As can be seen, the model can represent experimental data correctly. Figure 8 shows the composition evolution of samples with MnCl2 (1% w/w) exposed to 60% RH, the resulting adjusted model, with k1 ) 0.38, k2 ) 0.0031, and r ) 0.9941, and the residuals plot. Figure 9 shows the composition evolution of samples with 1%, 5%, 10%, and 20% FeCl3 (w/w) exposed to 80% RH and the resulting adjusted model. As can be seen in the figure, the model represents correctly the evolution of CaSO3 and CaSO4 concentrations, as well as the global product concentration (CaSO3 + CaSO4). From the model, by derivation with respect to time, it is possible to evaluate the formation rate (R) of calcium sulfite, calcium sulfate, or both of them (the global reaction rate):

R ) k1k2e-k2t When t ) 0, the initial reaction rate (R0) can be evaluated as R0 ) k1k2, with it possible to compare the effect of different foreign substances as well as the concentration of the added substance on the initial reaction rate. Figure 10 shows the dependence of resulting initial reaction rates (CaSO3 formation, CaSO4 formation, and CaSO3 + CaSO4 formation) for samples with different concentrations of FeCl3. As can be seen, the initial reaction rates show a linear dependence with the FeCl3 concentration. This dependence was observed in all of the samples studied with different concentrations of the same added substance. 4. Conclusions The effect of different substances on the reactivity of calcium carbonate exposed to sulfur dioxide has been studied. The main effect of those substances is the increase in the formation of calcium sulfite, the increase in the formation of calcium sulfate, or the increase in

Figure 10. Initial rate (R0) dependence with FeCl3 (1%, 5%, 10%, and 20% w/w) exposed to 80% RH (data arbitrarily displaced to avoid superposition).

the formation of both products with respect to pure calcium carbonate. In addition, the effect of RH has been studied. A mathematical model to represent the time evolution of the sample composition has been proposed and tested. The model is based on a capacity-limited heterogeneous reaction, and it represents correctly the evolution of the sample composition. The model allows one to compare quantitatively the effect of different substances added to calcium carbonate. In addition, it is possible to calculate reaction rates and to establish clearly the effect of the concentration of substances added to calcium carbonate on its reactivity. Literature Cited (1) Amorosso, G. G.; Fassina, V. Stone decay and conservationAtmospheric pollution, cleaning, consolidation and protection. Materials Science Monographs; Elsevier: Amsterdam, The Netherlands, 1983; Vol. 11. (2) Gauri, K. L.; Holdren, G. C. Pollutant effects on stone monuments. Environ. Sci. Technol. 1981, 15, 386. (3) Delgado Rodrıguez, J. Causes, mechanisms and measurement of damage in stone monuments. In Science, Technology and

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European Cultural Heritage; Baer, N. S., Sabbioni, C., Sors, A. I., Eds.; Butterworth-Heinemann: Oxford, 1991; pp 124-137. (4) Baedeker, P. A.; Reddy, M. M.; Reimann, K. J.; Sciammarella, C. A. Effects of acidic deposition on the erosion of carbonate stonesexperimental results from the U.S. national acid precipitation assesment programe. Atmos. Environ. 1992, 26B, 147. (5) Webb, A. H.; Bawden, R. J.; Busby, A. K.; Hopkins, J. N. Studies on the effects of air pollution on limestone degradation in Great Britain. Atmos. Environ. 1992, 26B, 165. (6) Camuffo, D.; del Monte, M.; Sabbioni, C. Origin and growth mechanisms of the sulfated crusts on urban limestone. Water, Air, Soil Pollut. 1983, 19, 351. (7) McGee, E. S.; Mossotti, V. G. Gypsum accumulation on carbonate stone. Atmos. Environ. 1992, 26B, 249. (8) Elfving, P. The initial steps in the air pollution induced deterioration of calcareous stone. In Proceedings of the 1995 LCP congress, Montreaux, 1995; Renato Pancella, Ed.; pp 421-427. (9) Gauri, K. L.; Popli, R.; Sarma, A. C. Effect of relative humidity and grain size on the reaction rates of marble at high concentrations of SO2. Durability Build. Mater. 1982, 1, 209. (10) Johansson, L. G.; Lindqvist, O.; Mangio, R. E. Corrosion of calcareous stones in humid air containing SO2 and NO2. Durability Build. Mater. 1988, 5, 439. (11) Spiker, E. C.; Comer, V. J.; Hosker, R. P.; Sherwood, S. I. Dry deposition of SO2 on limestone and marble: Role of humidity. In 7th International Congress on Deterioration and Conservation of Stone, Lisbon, 1992; Delgado Rodrigues, J., Henriques, F., Telmo Jeremias, F., Eds.; pp 397-406. (12) Pe´rez-Bernal, J. L. Estudio del sistema dio´xido de azufrecarbonato ca´lcico. Cata´lisis e inhibicio´n. Proteccio´n del patrimonio monumental. Ph.D. Dissertation, Universidad de Sevilla, Sevilla, Spain, 2000. (13) Kozlowski, R.; Hejda, A.; Ceckiewicz, S.; Haber, J. Influence of water contained in porous limestone on corrosion. Atmos. Environ. 1992, 26A, 3241. (14) Gauri, K. L.; Gwin, J. A. Deterioration of marble in air containing 5-10 ppm of SO2 and NO2. Durability Build. Mater. 1982, 1, 217. (15) Hoffmann, D.; Schimmelwitz, P.; Rooss, H. Interaction of sulfur dioxide with lime plasters. In Proceedings of the 2nd International Symposium on Deterioration of Building Stones, Athens, 1976; Beloyannis, N., Ed.; pp 37-42. (16) Haneef, S. J.; Jones, M. S.; Johnson, J. B.; Thompson, G. E.; Wood, G. C. Effects of air pollution on historic buildings and monuments 1986-1990. Scientific basis for conservation: Laboratory chamber studies. Eur. Cult. Heritage Newsletter Res. 1993, 7, 2. (17) Flatt, R.; Girardet, F.; Crovisier, J. L. Modelling of sulfur dioxide deposition on the Bern sandstone. In Proceedings of the 1995 LCP Congress, Montreaux, 1995; Renato Pancella, Ed.; pp 401-408. (18) Haneef, S. J.; Johnson, J. B.; Dickinson, C.; Thompson, G. E.; Wood, G. C. Effect of dry deposition of NOx and SO2 gaseous pollutants on the degradation of calcareous building stones. Atmos. Environ. 1992, 26A, 2963. (19) Radojevic, M.; Tyler, B. J.; Hall, S.; Penderghest, N. Air oxidation of sIV in cloud-water samples. Water, Air, Soil Pollut. 1995, 85, 1985.

(20) Hutchinson, A. J.; Johnson, J. B.; Thompson, G. E.; Wood, G. C.; Sage, P. W.; Cooke, M. J. The role of fly ash particulate matter and oxide catalysts in stone degradation. Atmos. Environ. 1992, 26A, 2795. (21) Ausset, P.; del Monte, M.; Lefevre, R. A. Embryonic sulfated black crusts on carbonate rocks in atmospheric simulation chamber and in the field-role of carbonaceous fly-ash. Atmos. Environ. 1999, 33, 1525. (22) Ausset, P.; Croivisier, J. L.; del Monte, M.; Furlan, V.; Girardet, F.; Hammecker, C.; Jannette, D.; Lefevre, R. A. Experimental study of limestone and sandstone sulphation in polluted realistic condition: The Lausanne atmospheric simulation chamber (LASC). Atmos. Environ. 1996, 30, 3197. (23) Sabbioni, C.; Zappia, G.; Gobbi, G. Carbonaceous particles and stone damage in a laboratory exposure system. J. Geophys. Res. 1996, 30, 3197. (24) Bo¨ke, H.; Go¨ktu¨rk, H.; Caner-Sa¨ltik, E.; Demirci, S. Effect of airborne particle on SO2-calcite reaction. Appl. Surf. Sci. 1999, 140, 70. (25) Vale Parapar, J.; Martın Pe´rez, A. Ensayo de Materiales en Atmo´ sferas Controladas; Servicio de Publicaciones de la Universidad de Sevilla: Sevilla, Spain, 1985. (26) Vale Parapar, J.; Martın Pe´rez, A. Ensayos de materiales en atmo´sferas simuladas. I. Criterios para el disen˜o de sistemas de simulacio´n. Mater. Constr. 1983, 189, 57. (27) Vale Parapar, J.; Martın Pe´rez, A. Ensayos de materiales en atmo´sferas simuladas. II. Disen˜o de un sistema de simulacio´n. Mater. Constr. 1983, 190, 53. (28) Vale Parapar, J.; Martın Pe´rez, A. Ensayos de materiales en atmo´sferas simuladas. III. Evaluacio´n de un sistema de simulacio´n. Mater. Constr. 1983, 192, 57. (29) Johnson, J. B.; Haneef, S. J.; Hepburn, B. J.; Hutchinson, A. J.; Thompson, G. E.; Wood, G. C. Laboratory exposure systems to simulate atmospheric degradation of building stone under dry and wet deposition conditions. Atmos. Environ. 1990, 24A, 2585. (30) Chun, K. C.; Quon, J. E. Capacity of ferric oxide particles to oxidize sulfur dioxide in air. Environ. Sci. Technol. 1973, 7, 532. (31) Ada´nez, J.; Gaya´n, P.; Garcıa-Labiano, F. Comparison of mechanistic models for the sulfation reaction in a broad range of particle sizes of sorbents. Ind. Eng. Chem. Res. 1996, 35, 2190. (32) Bravo, R. V.; Camacho, R. F.; Moya, V. M.; Garcıa, L. A. I. Desulphurization of SO2-NO2 mixtures by limestone slurries. Chem. Eng. Sci. 2002, 57, 2047. (33) Wilkins, S. J.; Compton, R. G.; Taylor, M. A.; Viles, H. A. Channel flow cell studies of the inhibiting action of gypsum on the dissolution kinetics of calcite: A laboratory approach with implications for field monitoring. J. Colloid Interface Sci. 2001, 236, 354. (34) Booth, J.; Hong, Q.; Compton, R. G.; Prout, K.; Payne, R. M. Gypsum overgrowths passivate calcite to acid attack. J. Colloid Interface Sci. 1997, 192, 207.

Received for review June 6, 2002 Revised manuscript received December 30, 2002 Accepted January 9, 2003 IE020426H