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Mechanisms of Metal Ion Sorption on Calcite: Composition Mapping by Lateral Force Microscopy Michael B. Hay,†,‡ Richard K. Workman,§ and Srinivas Manne*,† Department of Physics and Department of Materials Science & Engineering, University of Arizona, Tucson, Arizona 85721 Received July 17, 2002. In Final Form: November 4, 2002 We show that lateral force microscopy (also known as frictional force microscopy) can differentiate between substrate and overlayer phases during an inorganic surface reaction. A calcite substrate is imaged in situ, while immersed in aqueous solutions of pH ∼ 6-9 containing metal ions (Cd2+, Sr2+, and La3+) at concentrations of 10-5 to 10-3 M. Cd2+ and Sr2+ passivate surface steps, initiating overgrowth only in solutions already supersaturated relative to their respective carbonates. In contrast, La3+ initiates overgrowth even in undersaturated conditions and carries the reaction to completion by scavenging carbonate anions directly from the dissolving calcite surface. Monomolecular surface steps play a central role, serving as both dissolution sites for the substrate and nucleation sites for the overgrowth.
Introduction “Sorption” refers to the uptake of foreign ions by solid phases, whether it be by reversible adsorption to the surface, or chemical reaction with the surface, or eventual absorption into the bulk via solid-state diffusion.1 In environmental waters, ion-surface interactions can determine the geochemical fate of both the sorbate solution and the sorbent mineral. The concentrations of many trace metal ions in ocean waters, for instance, are several orders of magnitude lower than expected from geochemical weathering reactions.2 Simple precipitation cannot explain this scarcity, since metal ions in seawater are highly undersaturated with respect to their possible solid phases; therefore sorption and eventual incorporation of metal ions in existing minerals are implicated.1,2 Minerals are known to sequester radioactive ions,3,4 and sorption has often been suggested as a way to remove heavy metal pollutants from freshwater streams. Sorption has interesting and important effects on the mineral phase as well. For instance, the sorption of Mg2+ by calcite under growth conditions can induce a polymorphic transformation to aragonite,5 a form thermodynamically unfavorable under ambient conditions. Many marine organisms employ both calcite and aragonite biominerals in the construction of seashells,6 and biogenic control over the local ionic environment may be one way to achieve such phase transformations. Ion sorption therefore defines an important “inorganic baseline” for biomineralization. Apart from the specific geochemical relevance, ion sorption processes can also shed light on fundamental * Corresponding author. E-mail:
[email protected]. † Department of Physics. ‡ Current address: Department of Civil & Environmental Engineering, Princeton University, Princeton, NJ 08544. § Department of Materials Science & Engineering. (1) Recent review: Brown, G. E.; Parks, G. A. Intl. Geol. Rev. 2001, 43, 963-1073. (2) Krauskopf, K. B. Geochim. Cosmochim. Acta 1956, 10, 1-26. (3) Sturchio, N. C.; Antonio, M. R.; Soderholm, L.; Sutton, S. R.; Brannon, J. C. Science 1998, 281, 971-973. (4) Kersting, A. B.; Efurd, D. W.; Finnegan, D. L.; Rokop, D. J.; Smith, D. K.; Thompson, J. L. Nature 1999, 397, 56-59. (5) Kitano, Y.; Hood, D. W. J. Oceanogr. Soc. Jpn. 1962, 18, 35-39. (6) Lowenstam, H. A.; Weiner, S. On Biomineralization; Oxford University Press: New York, 1989.
mechanisms of surface reactions. When reactants interact via a dilute dispersed phase, the equilibrium composition of the reaction vessel can be predicted using solubility products. When one of the reactants is a solid, standard collision theory no longer holds,7 and barriers to reactant/ product mixing can kinetically stabilize the reaction far from thermodynamic equilibrium. Such transport barriers are highly sensitive to the specific reactant species and to the atomic order of the product phase. A classic example is metal passivation.8 When a surface of pure Al is exposed to oxygen, the surface quickly develops a passivating oxide “skin” that prevents O2 from reaching the underlying metal, stabilizing it from further corrosion. The extent of this effect varies widely among metals, and the molecular volume of the oxide is considered a critical factor in establishing an effective diffusion barrier. Scanning tunneling microscopy (STM) has recently shed light on such kinetic stabilization mechanisms,9,10 and some parallels between these studies and our current results are mentioned below. We describe the sorption mechanisms of Cd2+, Sr2+, and La3+ on the calcite cleavage plane. Calcite has been a substrate of choice for dissolution, growth, and sorption studies because it is naturally abundant, geochemically important, and easily cleaves into locally monomolecular planes with well-defined lattice structure and surface chemistry. Pioneering experimental contributions have come from the measurement of reaction rates by solution depletion methods and by elemental and/ or structural analysis of the solid product.11-29 More recently, several (7) Zaera, F. Acc. Chem. Res. 2002, 35, 129-136. (8) Adamson, A. W.; Gast, A. P. Physical Chemistry of Surfaces, 6th ed.; John Wiley & Sons: New York, 1997; pp 282-284. (9) Brune, H.; Wintterlin, J.; Trost, J.; Ertl, G.; Wiechers, J.; Behm, R. J. J. Chem. Phys. 1993, 99, 2128-2148. (10) Maurice, V.; Strehblow, H.-H.; Marcus, P. J. Electrochem. Soc. 1999, 146, 524-530. (11) Terjesen, S. G.; Erga, O.; Thorsen, G.; Ve, A. Chem. Eng. Sci. 1961, 14, 277-288. (12) Meyer, H. J. J. Cryst. Growth 1984, 66, 639-646. (13) Gutjahr, A.; Dabringhaus, H.; Lacmann, R. J. Cryst. Growth 1996, 158, 310-315. (14) Zachara, J. M.; Cowan, C. E.; Resch, C. T. Geochim. Cosmochim. Acta 1991, 55, 1549-1562. (15) Paquette, J.; Reeder, R. J. Geology 1990, 18, 1244-1247. (16) Paquette, J.; Reeder, R. J. Geochim. Cosmochim. Acta 1995, 59, 735-749.
10.1021/la020647s CCC: $25.00 © 2003 American Chemical Society Published on Web 03/20/2003
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groups have employed atomic force microscopy (AFM) to directly image and quantify calcite solution processes in situ at near-atomic length scales.30-42 Much of this work has focused on pure dissolution and growth in the absence of foreign ions.30-34,36-39 Dissolution occurs by an initial nucleation of monomolecular etch pits, followed by a steady retreat of faceted steps,30,39 consistent with the fast propagation of multiple kinks along each step site.38 Similarly, growth occurs by the steady advance of monomolecular steps, arising predominantly from spiral dislocations,31,33 but also (under certain solution conditions) from island nucleation.32,36 The presence of simple steps, islands, and single- and multiple-armed screw dislocations makes even the pure calcite cleavage plane reactively heterogeneous. A careful analysis by Dove et al.36 has shown that solution reaction rates measure a single composite value for a variety of simultaneous reaction sites, complicated further by the variety of crystal planes expressed in a typical powdered suspension; reaction rates therefore cannot uniquely identify local reaction mechanisms, and the constraints provided by AFM images lead to a clearer understanding of mineral surface reactions.36,43 Surface reactions of metal ions with calcite are inherently more complicated than pure calcite kinetics, because dissolution (of the calcite) and growth (of the new metal carbonate) can occur simultaneously, and the new phase itself may be either pure or a solid solution. Recent AFM (17) Reeder, R. J. Geochim. Cosmochim. Acta 1996, 60, 1543-1552. (18) Davis, J. A.; Fuller, C. C.; Cook, A. D. Geochim. Cosmochim. Acta 1987, 51, 1477-1490. (19) Stipp, S. L.; Hochella, M. F.; Parks, G. A.; Leckie, J. O. Geochim. Cosmochim. Acta 1992, 56, 1941-1954. (20) Chiarello, R. P.; Sturchio, N. C. Geochim. Cosmochim. Acta 1994, 58, 5633-5638. (21) Chiarello, R. P.; Sturchio, N. C.; Grace, J. D.; Geissbuhler, P.; Sorensen, L. B.; Cheng, L.; Xu, S. Geochim. Cosmochim. Acta 1997, 61, 1467-1474. (22) Prieto, M.; Fernandez-Gonzalez, A.; Putnis, A.; Fernandez-Dias, L. Geochim. Cosmochim. Acta 1997, 61, 3383-3397. (23) Mucci, A.; Morse, J. W. Geochim. Cosmochim. Acta 1983, 47, 217-233. (24) Pingitore, N. E.; Eastman, M. P. Geochim. Cosmochim. Acta 1986, 50, 2195-2203. (25) Pingitore, N. E.; Lytle, F. W.; Davies, B. M.; Eastman, M. P.; Eller, P. G.; Larson, E. M. Geochim. Cosmochim. Acta 1992, 56, 15311538. (26) Tesoriero, A. J.; Pankow, J. F. Geochim. Cosmochim. Acta 1996, 60, 1053-1063. (27) Sturchio, N. C.; Chiarello, R. P.; Cheng, L.; Lyman, P. F.; Bedzyk, M. J.; Qian, Y.; You, H.; Yee, D.; Geissbuhler, P.; Sorensen, L. B.; Liang, Y.; Baer, D. R. Geochim. Cosmochim. Acta 1997, 61, 251-263. (28) Cantrell, K. J.; Byrne, R. H. Geochim. Cosmochim. Acta 1987, 51, 597-605. (29) Zhong, S.; Mucci, A. Geochim. Cosmochim. Acta 1995, 59, 443453. (30) Hillner, P. E.; Gratz, A. J.; Manne, S.; Hansma, P. K. Geology 1992, 20, 359-362. (31) Gratz, A. J.; Hillner, P. E.; Hansma, P. K. Geochim. Cosmochim. Acta 1993, 57, 491-495. (32) Dove, P. M.; Hochella, M. F. Geochim. Cosmochim. Acta 1993, 57, 705-714. (33) Teng, H. H.; Dove, P. M.; Orme, C. A.; De Yoreo, J. J. Science 1998, 282, 724-727. (34) Teng, H. H.; Dove, P. M.; De Yoreo, J. J. Geochim. Cosmochim. Acta 1999, 63, 2507-2512. (35) Davis, K. J.; Dove, P. M.; De Yoreo, J. J. Science 2000, 290, 1134-1137. (36) Teng, H. H.; Dove, P. M.; De Yoreo, J. J. Geochim. Cosmochim. Acta 2000, 64, 2255-2266. (37) Park, N.-S.; Kim, M.-W.; Langford, S. C.; Dickinson, J. T. Langmuir 1996, 12, 4599-4604. (38) Liang, Y.; Baer, D. R.; McCoy, J. M.; LaFemina, J. P. J. Vac. Sci. Technol. A 1996, 14, 1368-1375. (39) Liang, Y.; Baer, D. R. Surf. Sci. 1997, 373, 275-287. (40) Astilleros, J. M.; Pina, C. M.; Fernandez-Diaz, L.; Putnis, A. Geochim. Cosmochim. Acta 2000, 64, 2965-2972. (41) Lea, A. S.; Amonette, J. E.; Baer, D. R.; Liang, Y.; Colton, N. G. Geochim. Cosmochim. Acta 2001, 65, 369-379. (42) Hoffmann, U.; Stipp, S. L. S. Geochim. Cosmochim. Acta 2001, 65, 4131-4139.
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work has explored the effects of Mg2+ and Ba2+ on calcite growth and the effects of Mn2+ and Sr2+ on calcite dissolution.35,40,41 Two of these papers35,41 intuit the surface reaction mechanism from the observed effects of the metal ions on step morphology and velocity; surface compositional variations due to sorption are not directly visualized. Only when the sorbate is Ba2+, an ion much (47%) larger than Ca2+, does a distinct sorbate phase become topographically visible against the adjoining calcite.40 In many cases of particular interest, however, sorbate ions (e.g., Cd2+) can be closely matched in size to Ca2+, making topographic distinctions between substrate and overgrowth very difficult. We show that lateral force microscopy or LFM, using frictional contrast between a substrate and a strained overlayer, can serve as a composition mapping technique during a surface reaction. LFM monitors the degree of cantilever twist using a four-quadrant photodiode, mapping this lateral force while the surface is scanned under constant normal force.44,45 Tip-sample friction depends on local surface chemistry and atomic structure, making LFM a sensitive probe of composition variations. Composition mapping has been used most often with organic films on inorganic substrates, where differences in hydrophobicity can yield large friction contrasts.46,47 However, LFM has shown distinct compositional contrast also in completely inorganic systems: Hydrogen-terminated Si vs SiO2,48 InP vs InGaAs,49 InAs and InSb on InP,50 C60 on NaCl,51 and MoS2 on mica.52 For those cases where a substrate and overlayer can clearly be identified,50-52 the overlayer always shows higher friction than the substrate, suggesting a common mechanism involving epitaxial strain in the overlayer. The same contrast is also observed in this work. While we limit our attention here to ion sorption overlayers, a contrast mechanism based on strainenhanced friction is likely quite general and may find future applications also in electrochemical or gas-phase reactions. Experimental Details Sorption experiments were performed with three cations (Cd2+, Sr2+, and La3+) whose affinities for CO32- exceed that of Ca2+. Table 1 shows the cation sizes and the structure and solubility of the corresponding thermodynamically stable carbonates. Among divalent cations, those smaller than Ca2+ (e.g., Cd2+) favor the rhombohedral calcite structure with metal ions six(43) The inability of adsorption isotherms to discriminate between competing molecular models is also illustrated in the interfacial selfassembly of surfactants. Here too adsorption isotherms have been interpreted with reference to an assumed molecular structure (typically flat monolayers and bilayers) that AFM results have shown to be too simplistic. See: Manne, S.; Gaub, H. E. Science 1995, 270, 1480-1482. Warr, G. G. Curr. Opin. Colloid Interface Sci. 2000, 5, 88-94. (44) Mate, C. M.; McClelland, G. M.; Erlandsson, R.; Chiang, S. Phys. Rev. Lett. 1987, 59, 1942-1945. (45) Recent review: Gnecco, E.; Bennewitz, R.; Meyer, E. J. Phys.: Condens. Mater. 2001, 13, R619-R642. (46) Overney, R. M.; Meyer, E.; Frommer, J.; Brodbeck, D.; Lu¨thi, R.; Howald, L.; Gu¨ntherodt, H.-J.; Fujihira, M.; Takano, H.; Gotoh, Y. Nature 1992, 359, 133-135. (47) Frisbie, C. D.; Rozsnyai, L. F.; Noy, A.; Wrighton, M. S.; Lieber, C. M. Science 1994, 265, 2071-2074. (48) Scandella, L.; Meyer, E.; Howald, L.; Lu¨thi, R.; Guggisberg, M.; Gobrecht, J.; Gu¨ntherodt, H.-J. J. Vac. Sci. Technol. B 1996, 14, 12551258. (49) Tamayo, J.; Gonzalez, L.; Gonzalez, Y.; Garcia, R. Appl. Phys. Lett. 1996, 68, 2297-2299. (50) Tamayo, J.; Garcia, R.; Utzmeier, T.; Briones, F. Phys. Rev. B 1997, 55, R13436-R13439. (51) Lu¨thi, R.; Meyer, E.; Haefke, H.; Howald, L.; Gutmannsbauer, W.; Gu¨ntherodt, H.-J. Science 1994, 266, 1979-1981. (52) Schumacher, A.; Kruse, N.; Prins, R.; Meyer, E.; Lu¨thi, R.; Howald, L.; Gu¨ntherodt, H.-J.; Scandella, L. J. Vac. Sci. Technol. B 1996, 14, 1264-1267.
Mechanisms of Metal Ion Sorption on Calcite
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Table 1. Structure and Solubility of Metal Carbonates
cation
ion radiusa (pm)
Ca2+
100
Cd2+
95
Sr2+
131
La3+
127
carbonate structure and lattice constantsb (nm) CaCO3 (calcite) hexagonal a ) 0.499, c ) 1.706 CdCO3 (otavite) hexagonal a ) 0.492, c ) 1.630 SrCO3 (strontianite) orthorhombic a ) 0.513, b ) 0.842, c ) 0.609 La2(CO3)3‚8H2O (lanthanite) orthorhombic a ) 0.898, b ) 0.958, c ) 1.700
solubility product,c Ksp 10-8.48 10-12.1 10-9.27 10-29.9
a Ion radii are taken from ref 70, using ion coordination numbers appropriate for each structure (6 for calcite and otavite, 9 for strontianite, and 10 for lanthanite). b Carbonate lattice data are taken from refs 63 (calcite and strontianite), 20 (otavite), and 71 (lanthanite). c Solubility products are taken from refs 72 (calcite), 73 (otavite), 74 (strontianite), and 65 (lanthanite).
coordinated to oxygens, whereas larger ones (e.g., Sr2+) favor the orthorhombic aragonite structure with metal ions nine-coordinated to oxygens. (Ca2+ itself can form both types of carbonates, although calcite is the thermodynamically favored form at ambient conditions.) The extent to which each cation species reacts with CaCO3 can be estimated using the individual Ksp values. For example, the net reaction for Cd2+ is
CaCO3 + Cd2+ f Ca2+ + CO32- + Cd2+ f Ca2+ + CdCO3 The net solubility product (using Table 1) is therefore
KCaCO3fCdCO3 )
Ksp,CaCO3 Ksp,CdCO3
≡
[Ca2+] [Cd2+]
) 4200
Therefore almost all of the initial Cd2+ reacts with molecular CaCO3, releasing Ca2+ in the process and driving the free [Ca2+] concentration much higher than the remaining [Cd2+]. As mentioned above, this analysis is only valid when the reactants are highly dispersed; the reaction with a solid cleavage plane of calcite is very different, as shown below. For each metal ion, the sorption mechanism was initially investigated in solutions undersaturated with respect to both calcite and the metal carbonate. Such “doubly undersaturated” conditions constrain the sorption pathway to either the direct adsorption of metal ions to the calcite surface or the cooperative dissolution of calcite and growth of metal carbonate. Any surface formation of metal carbonate must occur via calcite dissolution, because no other significant source of carbonate anion exists. By contrast, the reaction solutions in most previous work contained significant quantities of free CO32-, being either (i) supersaturated with respect to calcite but undersaturated with respect to the metal carbonate (the “coprecipitation” regime);12,23-25,29,35,40 or (ii) undersaturated with respect to calcite but supersaturated with respect to the metal carbonate (the “heterogeneous nucleation” regime, where the calcite acts as a substrate for the new growth);19-21 or (iii) supersaturated with respect to both phases.15-17,26 Investigations with doubly undersaturated solutions have been limited to sorption isotherms with powder samples11,13,14,18 and a recent AFM study41 that measured step speed effects but did not image surface reactions. In this work, solutions supersaturated with the metal carbonate were used only if doubly undersaturated solutions failed to produce an extensive surface reaction. Reaction solutions were simply the commercially available chloride salts of the sorbate ions (purity >99%) dissolved in water (distilled and deionized, resistivity > 18MΩ‚cm) to concentrations of 10-5 to 10-3 M. The free CO32- concentration, and hence the saturation state of the metal carbonate in solution, was controlled by either using the aqueous solution as is (pH ≈ 5.9) or by adjusting its pH to between 8 and 9 (using 0.1 M NaHCO3 and 0.1 M NaOH) and letting the solution come to equilibrium with
the ambient atmosphere.53 Specific concentrations and degrees of supersaturation are reported below in terms of the saturation index σ ≡ ln(K/Ksp), where K represents the appropriate product of actual ion concentrations in solution; thus σ < 0 for undersaturated and σ > 0 for supersaturated solutions.54 In each case, the calculation of K explicitly accounted for ion-pair association, using complexation constants found in the literature.14,28,55,56 We used a commercially available AFM/LFM system57 and monitored sample height and friction simultaneously, using a scan angle of 90° to minimize the mixing of the two signals. All images displayed in this work are unfiltered except for slope removal, unless otherwise indicated in the figure captions. Large scans of the surface typically used scan rates of 4-10 Hz and feedback to keep the tip-sample force constant in the 1-10 nN range. Friction and topography were displayed simultaneously, the latter in height and/or deflection form. Atomic level scans typically used scan rates of ∼20 Hz and were acquired nearly open loop (i.e., with low feedback gains) while monitoring the deflection and friction signals; lattice measurements were calibrated against the known calcite structure. While some overlayer/substrate friction contrast could be detected in every one of the experiments, the level of contrast varied unpredictably with the cantilever and imaging force. Usually, though not always, friction contrast was absent at very low forces and “turned on” abruptly at a finite force, perhaps as the tip penetrated a surface layer of hydrated ions. All images were acquired in the “trace” direction (left to right), where bright features in the LFM scans correspond to high friction. All metal carbonate overlayers showed higher friction than the calcite substrate, a contrast confirmed to be totally absent for control experiments of pure calcite growth. Optically clear and colorless “Iceland spar” calcite samples were cleaved with a razor blade (
