Fluoride Removal by Calcite: Evidence for Fluorite Precipitation and

Environ. Sci. Technol. 2005, 39, 9561-9568

Fluoride Removal by Calcite: Evidence for Fluorite Precipitation and Surface Adsorption BRETT D. TURNER,† P H I L I P B I N N I N G , * ,‡ A N D S . L . S . S T I P P § Civil, Surveying & Environmental Engineering, Newcastle University, University Drive, Callaghan, New South Wales 2308, Australia, Institute of Environment and Resources, Technical University of Denmark, Lyngby 2800, Denmark, and Geological Institut, University of Copenhagen, Øster Voldgade 10, DK-1350 Københaven K, Denmark

Fluoride contamination of groundwater, both anthropogenic and natural, is a major problem worldwide. In this study, fluoride removal by crushed limestone (99% pure calcite) was investigated by batch studies and surfacesensitive techniques from solutions with fluoride concentrations from 150 µmol/L (3 mg/L) to 110 mM (∼2100 mg/L). Surface-sensitive techniques, including atomic force microscopy (AFM) and X-ray photoelectron spectroscopy (XPS) as well as ζ potential measurements, confirm that, in addition to precipitation reactions, adsorption of fluoride also occurs. Results indicate that fluoride adsorption occurs immediately over the entire calcite surface with fluorite precipitating at step edges and kinks, where dissolved Ca2+ concentration is highest. The PHREEQ geochemical model was applied to the observed data and indicates that existing models, especially at low fluoride concentrations and high pH (>7.5) are not equipped to describe this complex system, largely because the PHREEQ model includes only precipitation reactions, whereas a combination of adsorption and precipitation parameters are required.

Introduction Until recently, fluoride was commonly thought to be removed from solution by precipitation in calcite systems. However, in 2003, Fan et al. (1) published a paper suggesting that sorption of fluoride onto calcite can occur. This paper extends that work by considering a much larger range of concentrations, examines the impact of pH on the system, and considers evidence for sorption when precipitation of fluorite is known to occur. We show that current models, considering only fluorite precipitation, cannot reproduce experimental results. Anthropogenic fluoride contamination of groundwater is associated with mineral processing industries including coalfired power stations, beryllium extraction plants, brick and iron works, and aluminum smelters. For example, at an aluminum smelter at Kurri Kurri, New South Wales, Australia, fluoride concentrations in groundwater range up to 160 mM (3000 mg/L). Elevated concentrations can also occur through the natural dissolution of fluoride-bearing minerals in bedrock and soil. In the northern area of the former Republic * Corresponding author phone: +45 4525 2161; e-mail: pjb@ er.dtu.dk. † Newcastle University. ‡ Technical University of Denmark. § University of Copenhagen. 10.1021/es0505090 CCC: $30.25 Published on Web 11/10/2005

 2005 American Chemical Society

of Bophuthatswana, South Africa, groundwater fluoride concentrations are as high as ∼3 mM (60 mg/L) (2). Boruff (3) showed that drinking water with fluoride concentrations above ∼150 µM causes dental and skeletal fluorosis, a mineralization disorder that impairs tooth and bone development. Many researchers (4-9) have used adsorption onto activated alumina or ion exchange (10) to remove fluoride from solution, while others (11-18) have used various forms of lime [Ca(OH)2, CaO, CaCO3] or other calcium salts to increase calcium activity in solution, thereby removing fluoride by precipitation as fluorite (CaF2). These techniques have advantages and disadvantages. For example, adsorption onto activated alumina requires pH < 7.0 before any real gains are made, but decreasing solution pH greatly increases the amount of aluminum species released into solution to potentially dangerous levels and frequent regeneration of sorption sites is required. The major disadvantages of precipitation methods are the difficulty in obtaining equilibrium fluoride concentrations less than ∼0.42 mM because of solubility constraints and pH above 9, which is unacceptable for some regulatory authorities. Calcite has been suggested as a fluoride removal medium by many authors. For example, Pickering and co-workers (15, 19, 20) concluded that fluorite precipitation was the main mechanism of fluoride removal because the mass of fluoride lost from solution was independent of the weight of calcite present. Reardon and Wang (11) employed data from column experiments to conclude that fluoride removal was achieved by dissolution of calcite and subsequent precipitation of CaF2. However, the hypothesis that precipitation of CaF2 is the only method of fluoride removal is inconsistent with the known behavior of calcite exposed to other ions. For example, calcite can also adsorb metal ions such as cadmium (21, 22), manganese (23), and zinc (24) as well as phosphate (25), with the degree of adsorption being pH-dependent. Yang et al. (13) expanded on work done by Glover and Sippel (14) and Simonsson (12) to include the effect of competing anions (SO42- and PO43-) on fluoride removal by calcite and found decreased fluoride removal in their presence. In addition, sulfate and phosphate were removed even when solutions of their respective calcium salts were undersaturated. Yang and colleagues could not explain their removal by precipitation but did not consider the possibility of adsorption. The moderate solubility of carbonate minerals, such as calcite, makes the application of adsorption theory difficult because of the problem of differentiating adsorption from precipitation and desorption from dissolution (26). To overcome this problem, Fan et al. (1) studied the adsorption kinetics of the 18F radioisotope onto calcite in solutions undersaturated with respect to fluorite. At pH 6.0 and an initial concentration of 0.57 µmol/L they found that fluoride adsorption was a pseudo-second-order process with the amount of fluoride adsorbed being approximately 3.6 × 1012 molecules/g of solid. However, Fan et al. considered only systems where fluoride concentrations are small and precipitation of fluorite does not occur. In this work, higher concentrations are considered and evidence for a combination of adsorption and precipitation mechanisms is presented. Results include data from traditional wet-chemical methods (batch tests) and observations from the surfacesensitive techniques atomic force microscopy (AFM) and X-ray photoelectron spectroscopy (XPS), as well as ζ-potential measurements. The evidence shows that adsorption can affect the ultimate mass of fluoride removed from solution. VOL. 39, NO. 24, 2005 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

9

9561

Materials and Methods Calcite. This study employed a crushed limestone from the Moore Creek & Sulcor Limestone member which is part of the Devonian (Frasnian) period Tamworth formation located near Tamworth, New South Wales, Australia. We used crushed limestone rather than commercial calcite powder to represent conditions in real water treatment systems. Analysis with a Philips PW1404 wavelength dispersive sequential XRF (X-ray fluorescence) analyzer, with less than (2% error for elemental analysis, showed that the limestone was 99% pure CaCO3 with trace quantities of Mg (1% as MgO), Si ( 0.97) are obtained, indicating that equilibrium fluoride concentration is directly proportional to pH. For the 150 µm fraction, for example, the percent removal of fluoride from a solution containing an initial concentration of 110 mM can be described by

% Fremoved ) -4.72pH + 133.77 Similar results were obtained by Yang et al. (13), who found that, theoretically, it is possible to decrease fluoride concentrations to any level by selecting the appropriate [H+]/[F-] ratio. The surface area dependence of fluoride removal can be observed in Figure 2. For an initial [F-]T ) 37 mM (Figure 2A), there is very little difference (Caδ+ and >CO3δ-, where > represents the edge of the bulk calcite surface). This charge imbalance is decreased by the hydrolysis of water, resulting in the formation of >Ca-OHδ- and >CO3-Hδ+ species. The formation of these species is thermodynamically favorable as the added hydrolysis layer lowers the potential with respect to the bulk calcite surface. These species can also react with any other ions in solution. For example, decreasing the system pH (increasing H+ and Ca2+) causes an increase in positive sorption via the reactions:

>Ca-OHδ- + H+ T >Ca-OH2δ+ T >Caδ+ + Η2Ο >CO3Hδ+ + Ca2+ T >CO3Caδ+ + Η+ At neutral to higher pH, OH- and CO32- dominate, producing >CaCO3- and >CO3+ sorption sites (44). Therefore, as well as providing more Ca2+ in solution from the dissolution of calcite, decreasing pH effectively increases the number of positive surface sites leading to an increase in the amount of fluoride removed from the system. Changes in the surface potential of particles can be followed by ζ potential measurements. ζ Potential Measurements. Fluoride adsorption onto the surface contributes negative charge, so it should be expected that adsorption causes a shift in the point of zero charge, pHzpc, to a lower value. Experiments investigating the effect of F- adsorption on the surface charge of calcite show that with increasing F- concentration the pHzpc shifts from ∼6.3 to ∼2.6 (see Supporting Information, Figure S2) indicating F adsorption at the calcite surface. X-ray Photoelectron Spectroscopy. XPS can be used to examine the calcite surface in detail, providing data on the composition and form of surface species. XPS spectra (Supporting Information, Figure S3) were obtained for material from a batch experiment with 1.18 mm size calcite reacted for 3 days with a solution that originally contained 35 mM NaF, together with a similar sample exposed only to DI water for the same amount of time. On the sample exposed to NaF, strong F peaks are observed near 685 and 830 eV, indicating that significant F was taken up by the surface. These peaks were not found in the sample exposed only to DI water. Thus, the F that we observe associated with the surface was either adsorbed or precipitated or both during the period of the samples’ exposure to the NaF solution. The peak intensity ratios (Supporting Information Table S2) show that the sample exposed to NaF solution for 3 days had a F/Ca ratio between 1 and 2 and a F/C ratio suggesting roughly equivalent concentrations of F and CO3 in the nearsurface. We can therefore conclude that the fluorite formed as separate crystallites with F adsorbed over calcite in the spaces between. Atomic Force Microscopy. Atomic force microscopy was used to observe changes in morphology as a result of fluoride exposure and thereby to determine the mechanisms involved in its uptake. Examination of surface behavior with AFM requires a sample that is flat on the micrometer scale. The crushed limestone that was used in the batch experiments is very fine-grained, and although fractured surfaces appeared flat under an optical microscope, their roughness stretched the capabilities of the AFM scanner and cantilever to their limits. Freshly broken fragments did provide images where the surfaces could occasionally be seen through the noise. We were able to observe atomically flat {101h 4} terraces with widths of tenths of a micrometer to a micrometer or two, often separated by steps of one or two atomic layers. On samples exposed to deionized water, rhombohedral etch pits formed. Under both conditions, surfaces behaved exactly as expected from numerous, previous investigations of cleaved

FIGURE 3. Height mode AFM images of the {101h4} surface of calcite: (A) exposed for several minutes only to water, showing smooth terraces and typical calcite etch pit formation; (B) exposed for 10 min to 5 mM NaF, showing roughened terraces and atypical etch pit formation. In height mode images, features close to the observer are light in color, while those farther away are dark. calcite specimens where crystal size was 2 mm across or larger (45, 46). This similarity of behavior justified working with cleaved calcite specimens, which gave clearer data and allowed for more complicated AFM experiments. Investigations focused on the first minutes following exposure and then ensured consistency by checking samples exposed for several days. Figure 3A shows an image of the control, a sample of calcite that was exposed only to deionized water for several minutes and dried by physically removing the remaining solution with a N2 jet. Terraces are broad, atomically flat, and separated by steps of 3 Å, the thickness of a molecular layer. Etch pits are regular and rhombohedral, following the preferred cleavage and growth directions for calcite. Typically, etch pits expand along each atomic layer until they reach several micrometers in width before a new pit nucleates and enlarges in the level below. Dissolution removes material from step edges and from defect sites in the terraces. Figure 3B shows a fresh sample exposed to 5 mM NaF for 10 min and then removed from solution and dried in order VOL. 39, NO. 24, 2005 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

9

9565

to halt surface processes so the effects could be examined. Although pH is near 8, the solution is undersaturated with respect to calcite, so dissolution erodes the surface. Steps are still one atomic layer high, and step edges still retreat one layer at a time, but now pits are elliptical and terrace edges are irregular with long, pointed peninsulas. Terraces after exposure to NaF are no longer atomically flat and smooth, demonstrating that material has been removed to create a multitude of tiny pits, one atomic layer deep and 10-20 nm across. We observe this in Figure 3B as graininess in the terraces. Experiments with solutions of Na2CO3, H2CO3, CaCO3, and HNO3 (47) reveal dissolution behavior very similar to that observed in pure water (Figure 3A), so we interpret that the presence of F-, not Na+, is responsible for the very rough and irregular morphology we see in Figure 3B. The system is still driven to dissolve by undersaturation with respect to Ca2+ and CO32-, but dissolution in the presence of F- is both enhanced on terraces to produce holes and blocked at some sites on step edges to produce atypical morphology. This behavior can only be explained by Fadsorption over the entire surface. Figure 4A shows a similar sample exposed to 35 mM NaF for 10 min and then removed from solution and dried. Traces of elliptical etch pits and rough terrace surfaces are again visible (top right, arrow), but two sizes of crystallites have precipitated. Some small, rounded crystallites are scattered randomly over terraces, and some sit as ordered decorations along terrace. In other locations, huge cubic crystals modify the retreat of calcite steps around them (lower right, arrows). Such formation of crystallites is not observed on the control samples or on surfaces exposed to supersaturated solutions of CaCO3 (47). Longer exposure to NaF solution or exposure to solutions of higher F concentration results in larger crystallites. This, together with results from the batch experiments and XPS study, lead to the conclusion that the crystallites are CaF2. In the AFM images, we see fluorite nucleate and grow very quickly, within seconds, leading us to conclude that growth is limited by the provision of Ca2+ as calcite dissolves. Thus, fluorite precipitation at or near step edges is logical because Ca2+ concentration in this area is highest as a result of step retreat. Erosion of calcite continues after nucleation and initial growth of the crystallites, as we see in the curved step edges to the right of the cubic crystals in Figure 4A and also in Figure 4B. This last image was taken from a sample exposed to 5 mM NaF for only 1 min. Even after such a short time in a relatively dilute solution, adsorption and dissolution have already roughened surfaces, small crystallites have precipitated on terraces and near step edges, and these edges have retreated beyond the site of initial crystallite nucleation, providing material for further crystal growth. The arrows on Figure 4B are vectors, pointing in the direction of edge retreat during crystallite growth. Samples exposed to solution for several days, simulating conditions of the batch experiments, were rougher, with larger crystallites, but overall behavior was consistent. The AFM observations, together with the results from the batch and ζ potential studies, allow us to construct a conceptual model of the processes that take place at the surface. From the instant the calcite comes into contact with the NaF solution, F- adsorbs and calcite dissolves. We know adsorption is instantaneous because surface morphology is immediately affected. Dissolution increases Ca2+ concentration in solution, saturation with respect to CaF2 is reached, and fluorite precipitates where local concentration gradients are highest, at the edges of retreating terraces. Dissolution continues, despite adsorbed F- and precipitating fluorite, replenishing dissolved Ca2+, allowing further precipitation of CaF2, and the cycle continues. As long as calcite can dissolve and supply Ca2+, fluorite continues to precipitate. 9566

9

ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 39, NO. 24, 2005

FIGURE 4. Deflection mode AFM images of calcite: (A) exposed for 10 min to 35 mM NaF; (B) exposed for 1 min to 5 mM NaF. Deflection mode images enhance morphology by highlighting slope. They are particularly useful for very rough surfaces. On these images, negative slopes are light and positive are dark, making it appear that a light is shining from the right side. The large features in panel A are cubes, as can be seen by the square appearance of the top of the largest. The sloping edges are an artifact: the tip’s pyramidal shoulders prevent its point from reaching the base of the cube. Smaller cubes (left) appear simply as a reflection of the pyramidal tip. The arrows of the right image are vectors that point in the direction of step retreat. Implications. This paper extends current knowledge of the geochemistry of fluoride removal by calcite. It has been shown that a combination of surface adsorption and precipitation reactions remove fluoride from aqueous systems; the degree of removal is dependent on calcite surface area. Calcite/limestone has been used for many years for the removal of fluoride from water, but the geochemical processes are much more complex than previously thought. Because geochemical models such as PHREEQ only include simple fluorite precipitation, they are unable to describe the observed results. A parameter describing adsorption is necessary. Thus, the process engineer should bear in mind that particle size does matter and should not rely only on geochemical software packages to predict the outcome of the fluoride/limestone treatment process.

Acknowledgments Our thanks go to Dieke Postma of the Danish Technical University (DTU) for his comments. XPS analysis was kindly performed by the Department of Physics Surface Science research group at the University of Newcastle, New South Wales, Australia. This research was funded by Hydro Aluminium Kurri Kurri, the University of Newcastle, the Australian Research Council, the Danish Research Council, and the Carlsberg Foundation.

Supporting Information Available Detailed experimental methods, including batch test, pH, fluoride and calcium analyses, ζ potential measurements, and X-ray photoelectron spectroscopy. Additional results and discussion, including surface area and Ca atom surface site density, pH and surface area with data from low concentration fluoride batch tests on 150 µm calcite, ζ potential measurements, and X-ray photoelectron spectroscopy. This material is available free of charge via the Internet at http://pubs.acs.org.

Literature Cited (1) Fan, X.; Parker, D. J.; Smith, M. D. Adsorption kinetics of fluoride on low cost materials. Water Res. 2003, 37, 4929-4937. (2) McCaffrey, L. P. Distribution and causes of high fluoride groundwater in the Western Bushveld area of South Africa. Ph.D. Thesis, University of Cape Town, Cape Town, South Africa, 1998. (3) Boruff, C. S. Removal of fluoride from drinking waters. Ind. Eng. Chem. 1934, 26(1), 69-71. (4) Bishop, P. L. Fluoride removal from drinking water by fluidized activated alumina adsorption. J.sAm. Water Works Assoc. 1978, 554-559. (5) Clifford, D. Activated alumina: rediscovered “adsorbent” for fluorine, humic acids and silica. Ind. Water Eng. 1978, 15. (6) Hao, O. J. Adsorption characteristics of fluoride onto hydrous alumina. J. Environ. Eng. 1986, 112, 1054-1069. (7) Meenakshi, S.; Pius, A.; Karthikeyan, G.; Appa Rao, B. V. The pH dependence of efficiency of activated alumina in defluoridation of water. Indian J. Environ. Prot. 1991, 11, 511-513. (8) Rubel, F. J. The removal of excess fluoride from drinking water by activated alumina. J.sAm. Water Works Assoc. 1979, (January), 45-49. (9) Schoeman, J. J.; Botha, G. R. Evaluation of activated alumina process for fluoride removal from drinking water and some factors influencing its performance. Water SA 1985, 11, 25-32. (10) Singh, G. K. B.; Majumdar, J. Removal of fluoride from spent pot liner leachate using ion exchange. Water Environ. Res. 1999, 71(1), 36-42. (11) Reardon, E. J.; Wang, Y. A limestone reactor for fluoride removal from wastewaters. Environ. Sci. Technol. 2000, 34, 3247-3253. (12) Simonsson, D. Reduction of fluoride by reaction with limestone particles in a fixed bed. Ind. Eng. Chem. Process Des. Dev. 1979, 18, 288-292. (13) Yang, M.; Hashimoto, T.; Hoshi, N.; Myoga, H. Fluoride removal in a fixed bed packed with granular calcite. Water Res. 1999, 33, 3395-3402. (14) Glover, E. D.; Sippel, R. F. Experimental pseudomorphs: replacement of calcite by fluorite. Am. Mineral. 1962, 47, 11561165. (15) Farrah, H.; Slavek, J.; Pickering, W. F. Fluoride sorption by soil components: calcium carbonate, humic acid, manganese dioxide and silica. Aust. J. Soil Res. 1985, 23, 429-439. (16) Saha, S. Treatment of aqueous effluent for fluoride removal. Water Res. 1993, 27, 1347-1350. (17) Phantumvanit, P.; Legeros, R. Z. Characteristics of bone char related to efficacy of fluoride removal from highly fluoridated water. Fluoride 1997, 30, 207-218. (18) Wadhwani, T. K. The Mechanism of fluorine removal by calcium salts - part 1. Indian Inst. Sci. 1954. (19) Slavek, J.; Farrah, H.; Pickering, W. F. Interaction of clays with dilute fluoride solutions. Water, Air, Soil Pollut. 1984, 23, 209220. (20) Pickering, W. F. The mobility of soluble fluoride in soils. Environ. Pollut. (Ser. B) 1985, 9, 281-308.

(21) Davis, J. A.; Fuller, C. C.; Cook, A. D. A model for trace metal sorption processes at the calcite surface: adsorption of Cd2+and subsequent solid solution formation. Geochim. Cosmochim. Acta 1987, 51, 1477-1490. (22) van der Weijden, R. D.; Meima, J.; Comans, R. N. J. Sorption and sorption reversibility of cadmium on calcite in the presence of phosphate and sulfate. Mar. Chem. 1997, 57, 119-132. (23) McBride, M. B. Chemisorption and precipitation of Mn2+ at CaCO3 surfaces. Soil Soc. Am. J. 1979, 43, 693-698. (24) Zachara, J. M.; Kittrick, J. A.; Harsh, J. B. The mechanism of Zn2+ adsorption on calcite. Geochim. Cosmochim. Acta 1988, 52, 2281-2291. (25) Freeman, J. S.; Rowell, D. L. The adsorption and precipitation of phosphate onto calcite. J. Soil Sci. 1981, 32(1), 75-84. (26) Stipp, S. L. S. Toward a conceptual model of the calcite surface: hydration, hydrolysis, and surface potential. Geochim. Cosmochim. Acta 1999, 63, 3121-3131. (27) Gregg, S. J.; Sing, K. S. W. Adsorption, surface area and porosity, 2nd ed.; Academic Press: London, 1982. (28) Eaton, A. D.; Clesceri, L. S.; Greenberg, A. E. Standard Methods for the Examination of Water and Wastewater, 19th ed.; American Public Health Association: Washington, DC, 1995. (29) Moulder, J. F.; Stickle, W. F.; Sobol, P. E.; Bomben, K. D. Handbook of X-ray Photoelectron Spectroscopy; Physical Electronics: Eden Prairie, MN, 1995. (30) Stipp, S. L. S.; Hochella, M. F. J. Structure and bonding environments at the calcite surface as observed with X-ray Photoelectron Spectroscopy (XPS) and Low Energy Electron Diffraction (LEED). Geochim. Cosmochim. Acta 1991, 55, 17231736. (31) Parkhurst, D. L.; Appelo, C. A. J. User’s Guide to PHREEQC (Version 2)sA computer program for speciation, batch reaction, one-dimensional transport, and inverse geochemical calculations; U.S. Geological Survey Water-Resources Investigations Report 99-4259; USGS: Denver, CO, 1999. (32) Bond, A. M.; Hefter, G. T. Critical Survey of Stability Constants and Related Thermodynamic Data of Fluoride Complexes in Aqueous Solution; Pergamon Press: Oxford, U.K., 1980. (33) Ball, J. W.; Nordstrom, D. K. WATEQ3F-user’s manual with revised thermodynamic database and test cases for calculating speciation of major, trace and redox elements in natural waters; U.S. Geological Survey Open-File Report; USGS: Denver, CO, 1991. (34) Allison, J. D.; Brown, D. S.; Novo-Gradac, K. J. MINTEQA2/ PRODEFA2sA geochemical assessment model for environmental systemssVersion 3.0 User’s Manual; Office of Research and Development, U. S. Environmental Protection Agency: Athens, GA, 1990. (35) Appelo, C. A. J.; Postma, D. Geochemistry, Groundwater and Pollution; A. A. Balkema: Rotterdam, The Netherlands, 1999. (36) Tan, K. H. Principles of Soil Chemistry; Marcel Dekker: New York, 1982. (37) Siffert, B.; Fimbel, P. Parameters affecting the sign and the magnitude of the electrokinetic potential of calcite. Colloids Surf. 1984, 11, 377-389. (38) Foxall, T.; Peterson, G. C.; Rendall, H. M.; Smith, A. L. Charge determination at calcium salt/aqueous solution interface. J. Chem. Soc., Faraday Trans. 1979, 75(1), 1034-1039. (39) Somasundaran, P.; Agar, G. E. The zero point of charge of calcite. J. Colloid Interface Sci. 1967, 24, 433-440. (40) Thompson, D. W.; Pownall, P. G. Surface electrical properties of calcite. J. Colloid Interface Sci. 1988, 131, 74-82. (41) Huang, Y. C.; Fowkses, F. M.; Lloyd, T. B.; Sanders, N. D. Adsorption of calcium ions from calcium chloride solutions onto calcium carbonate particles. Langmuir 1991, 7, 17421748. (42) Cicerone, D. S.; Regazzoni, A. E.; Blesa, M. A. Electrokinetic properties of the calcite/water interface in the presence of magnesium and organic matter. J. Colloid Interface Sci. 1992, 154, 423-433. (43) van Cappellen, P.; Charlet, L.; W., S.; Wersin, P. A surface complexation model of the carbonate mineral-aqueous solution interface. Geochim. Cosmochim. Acta 1993, 57, 3505-3518. (44) Fenter, P.; Geissbuhler, P.; Srajer, G.; Sorenson, L. B.; Sturchio, N. C. Surface speciation of calcite observed in situ by highresolution X-ray reflectivity. Geochim. Cosmochim. Acta 2000, 64, 1221-1228. VOL. 39, NO. 24, 2005 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

9

9567

(45) Stipp, S. L. S.; Gutmannsbauer, W.; Lehmann, T. The dynamic nature of calcite surfaces in air. Am. Mineral. 1996, 81, 1-8. (46) Stipp, S. L. S.; Konnerup-Madsen, J.; Franzreb, K.; Kulik, A.; Mathieu, H. J. Spontaneous movement of ions through calcite at standard temperature and pressure. Nature 1998, 396, 356359. (47) Christensen, J. T. De sjældne jordarter i naturlig calcit og syntetisk Eu-CaCO3 som model for actinid optagelse. (The rare Earth

9568

9

ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 39, NO. 24, 2005

elements in natural calcite and synthetic Eu-CaCO3 as a model for actinide uptake).MSc Thesis, University of Copenhagen, Denmark, 2004.

Received for review March 16, 2005. Revised manuscript received October 4, 2005. Accepted October 6, 2005. ES0505090