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Soil Science Society of America Journal 64:1968-1975 (2000)
© 2000 Soil Science Society of America

DIVISION S-2-SOIL CHEMISTRY

Adsorption of Mercury(II) by Kaolinite

D. Sarkara, M.E. Essingtonb and K.C. Misrac

a Earth and Physical Sci. Div., Univ. of Texas, San Antonio, TX 78249-0663 USA
b Dep. of Plant and Soil Sci., The Univ. of Tennessee, Knoxville, TN 37901-1071 USA
c Dep. of Geological Sci., The Univ. of Tennessee, Knoxville, TN 37996-1410. Contribution from the Dep. of Plant and Soil Sciences, The Univ. of Tennessee USA

messington{at}utk.edu


   ABSTRACT
TOP
 ABSTRACT
INTRODUCTION
 Materials and methods
 Results and discussion
 Conclusions
 REFERENCES
 
Adsorption of Hg(II) by kaolinite was investigated as a function of solution pH, ionic strength, and the competitive or complexation effects of ligands (Cl, SO4, PO4) and metals (Ni and Pb). Mercury(II) adsorption from a 0.6 µM Hg(II) solution was primarily influenced by pH. The Hg(II) adsorption edge was described by a pH50 (pH where 50% adsorption occurs) of 3.4 and a pHmax (pH where maximum adsorption occurs) of 4.4. At pH values above the pHmax, Hg(II) retention decreased with increasing pH. Chloride and Ni shifted pH50 from 3.4 to 7 and 4.1, respectively. Nickel and Pb reduced the amount of Hg(II) adsorbed throughout the pH range examined. Ionic strength and the presence of SO4 and PO4 had relatively little impact on the Hg(II) adsorption envelope. The adsorption of Hg(II) was predicted through the application of the triple layer model (TLM) by assuming that the kaolinite surface was composed of equal proportions of silanol and aluminol groups. The TLM model suggests that the silanol group was responsible for retaining the bulk of the adsorbed Hg(II), through the formation of the {equiv}SiO--HgOH+ outer-sphere, and the {equiv}SiOHg -2 and {equiv}SiOHgCl0 or {equiv}SiOHgOHCl- (Cl system) inner-sphere species. The {equiv}AlO--HgOH+ outer-sphere complex accounted for a small percentage (<15–35%) of the adsorbed Hg(II). The TLM results suggested that Hg(II) adsorption by both {equiv}SiOH and {equiv}AlOH sites on kaolinite should be considered to predict adequately Hg(II) retention.

Abbreviations: IS, ionic strength • TLM, triple layer model


   INTRODUCTION
TOP
 ABSTRACT
 INTRODUCTION
Materials and methods
 Results and discussion
 Conclusions
 REFERENCES
 
MERCURY(II) [Hg(II)] enters the environment from six major sources: production and consumption of goods containing Hg(II); materials processing; mining and smelting; fuel burning; final disposal of waste streams; and natural sources (Watson, 1979). In all cases, Hg(II) may attain environmentally significant concentrations and pose serious health hazards. Like most other trace elements, the fate of Hg(II) in natural aquatic systems is largely controlled by adsorption reactions with fixed or mobile adsorbents (Schuster, 1991). In addition to affecting the physical transport of dissolved species, adsorption can lead to changes in chemical reactivity and biological activity. The clay minerals are considered very important adsorbents in natural water systems because of their high specific surface area combined with the structural- and pH-dependent charge developed on their surfaces. Kaolinite is a very common clay mineral in acidic soils and in weathered feldspathic aquifers. A detailed, quantitative understanding of Hg(II) adsorption on kaolinite is therefore important in determining the ultimate fate of this element in such environments.

Adsorption is a complex process involving physical, chemical, and electrical interactions at sorbent surfaces which can be exceedingly more complicated than reactions in the bulk solution (Riese, 1982). As a result, several conceptual and mathematical site-binding models (e.g., Davis and Leckie, 1978; Bowden et al., 1973; Farley et al., 1985) have been proposed to explain and predict metal adsorption behavior. These models, usually termed surface complexation models, require a knowledge of the properties of the oxide and oxide-like surfaces and assumptions about the location and form of the sorbed ions in the electrical double layer near the surface. Surface complexation models have been widely and successfully used to describe the adsorption of ions on simple systems involving pure oxides (Davis and Leckie, 1978; Westall and Hohl, 1980; Hayes and Leckie, 1987; Cowan et al., 1991); however, as Cowan et al. (1991) points out, their applicability to complex systems, where elements in solution may compete with each other for sorption sites, has received only limited attention. Moreover, our ability to predict quantitatively ion adsorption on multi-component oxides is still limited (Honeyman, 1984; Anderson and Benjamin, 1990). Kaolinite is one such multi-component oxide system, containing silanol ({equiv}SiOH) and aluminol ({equiv}AlOH) surface functional groups. Because of its complex surface chemistry, few investigators have quantitatively linked the measured surface properties of kaolinite with its ion adsorption capacity (Riese, 1982; Liechti, 1983; Zachara et al., 1988; Ward and Bassett, 1990; He et al., 1997). In all these studies, ion adsorption on the kaolinite surface has been successfully modeled as a composite of {equiv}SiOH and {equiv}AlOH sites, which is conceptually quite similar to the Patch Site Distribution (PSD) model of Meng and Letterman (1993) for modeling ion adsorption on Al-hydroxide modified silica.

The objectives of the present investigation were to: (i) examine adsorption of Hg(II) by kaolinite as a function of solution pH, ionic strength, and the presence of metals (Pb and Ni) and ligands (Cl, SO4, and PO4); and (ii) apply the triple layer surface complexation model to the experimental data and evaluate the ability of the model to predict Hg(II) adsorption by assuming the kaolinite surface to be a composite of {equiv}SiOH and {equiv}AlOH functional groups.


   Materials and methods
TOP
 ABSTRACT
 INTRODUCTION
 Materials and methods
Results and discussion
 Conclusions
 REFERENCES
 
Preparation of Solid
Kaolinite was prepared following the procedure of Mattigod et al. (1985). A well crystallized Georgia kaolinite (KGa-1, from the Source Clays Repository of the Clay Minerals Society) was used for the adsorption experiments. A distilled water suspension of the clay was dispersed in a blender for approximately 45 min. The suspension was then adjusted to pH 9.5 with 0.1 M NaOH, and the 0.2- to 2.0-µm size fraction was obtained by centrifugation. The size-separate was rinsed with 1 M NaNO3 and adjusted to pH 3 with HNO3. Following this treatment, the solid was centrifuged and rinsed with 0.01 M NaNO3 several times until the pH of the supernatant was approximately 6. The solid was then stored wet under N2 until needed.

Preparation of Solutions
All reagents were of analytical grade or better. A solution of 2.94 mM Hg(II) was prepared from Hg(NO3)2·H2O by dissolving the salt in a matrix of 10% (v/v) HNO3. The stock was diluted to prepare working solutions of 2.94 µM Hg(II) which underwent a further five-fold dilution in reaction vessels to yield a final Hg(II) concentration of 0.6 µM. The background electrolyte solutions were 0.01, 0.1, and 0.5 M NaNO3. Stock solutions of Pb and Ni were prepared from their nitrate salts and were diluted such that their final concentrations in the reaction vessels were 14 and 48 µM, respectively. Stock solutions of the ligands, Cl, SO4, and PO4 were prepared from their sodium salts with appropriate dilutions to yield final concentration of 0.01 M in the reaction vessels. The pH was controlled using 0.02 M NaOH and 0.02 M HNO3.

Analytical Techniques
Mercury(II) concentrations were determined by a Perkin Elmer Flow Injection Mercury System Atomic Absorption Spectrometer, model FIMS-400 (Perkin Elmer Corp., Norwalk, CT). Dilution of the samples prior to analysis was necessary and was accomplished with 10% HNO3 spiked with a few drops of 5% (w/v) KMnO4 to stabilize Hg(II). A Perkin Elmer PE-5000 Atomic Absorption Spectrophotometer was used for the determination of Pb and Ni. A Dionex Ion Chromatograph (model DX-100, Dionex Corp., Sunnyvale, CA) was used to determine the concentrations of the ligands. The pH of all solutions was measured using a combination glass pH electrode and an Orion EA 940 Ion Analyzer (Orion Research, Boston, MA).

Experimental Techniques
Mercury(II) Adsorption Experiments
The experimental procedure involved the equilibration of 0.1 g kaolinite in 24 mL of the electrolyte solution and 6 mL of the Hg(II) working solution in capped 50-mL Teflon centrifuge tubes. This resulted in concentrations of 3.3 g L-1 kaolinite and 0.6 µM Hg(II). Solution pH was adjusted with 0.02 M HNO3 or 0.02 M NaOH, such that the equilibrium solutions had pH values ranging from 3 to 9. Preliminary kinetic studies indicated that adsorption is characterized by a rapid initial uptake followed by a much slower, continuous uptake. A 48-h reaction period was found to be sufficient to achieve equilibrium conditions when using a reciprocating floor-shaker for continuous agitation. Triplicate equilibrations were performed at ambient temperature (20–25°C). Following equilibration, the solid and liquid phases were separated by centrifugation at 1500 x g for 40 minutes. Filtration was not used because of the high loss of Hg(II) by adsorption on the 0.45-µm membrane filter and its support. After centrifugation, a 3-mL aliquot of the supernatant was withdrawn by pipette for Hg(II) analysis. During preliminary evaluations, the adsorption tube was emptied and rinsed with deionized water and refilled with 30 mL of 10% HNO3. This tube was again shaken for 24 h and an aliquot of the acid rinse was withdrawn for analysis. This rinse step was used to account for any Hg(II) retained by the Teflon centrifuge tubes. Results showed that the centrifuge tube walls did not retain Hg(II). Mercury(II) adsorbed by kaolinite was calculated from the difference between the Hg(II) initially added to the system and that remaining in the solution after equilibration. The dilutions induced by the pH controls were considered while computing the amount of Hg(II) adsorbed. All experiments were conducted under N2 to eliminate the potential for precipitation of Hg(II) carbonate.

Binary Mercury(II)-Metal and Mercury(II)-Ligand Adsorption Experiments
The adsorption of Hg(II) in the presence of metals (Ni and Pb) and ligands (Cl, SO4, and PO4) was examined by batch experiments performed in triplicate and in 50-mL Teflon centrifuge tubes at ambient temperatures. For the experiments involving Hg(II) adsorption in the presence of Ni or Pb, the kaolinite suspension was prepared by adding 21 mL of 0.1 M NaNO3 to 0.1 g kaolinite, followed by the addition of 3 mL of stock solution of the metal such that the initial concentrations of the metals were 14 µM Pb and 48 µM Ni. A 6-mL aliquot of the Hg(II) working solution containing 2.94 µM Hg(II) was then added to make the total volume of the suspension 30 mL. The kaolinite suspensions spiked with metals were not allowed to equilibrate before the addition of Hg(II). For Hg(II) adsorption in the presence of ligands, 3 mL of stock ligand solution was added to the 0.1 g kaolinite suspended in 21 mL of 0.1 M NaNO3 and 6 mL of Hg(II) working solution. The concentrations of Cl, SO4, and PO4 were 0.01 M. The pH of the suspensions was adjusted with 0.02 M NaOH and 0.02 M HNO3, such that the equilibrium solutions had pH values ranging between 3 and 9. Following the 48-h equilibration period, the tubes were centrifuged at 1500 x g for 40 minutes and the supernatant solutions analyzed for both Hg(II) and the metal or ligand present in the system.

Adsorption Modeling
The TLM was used to describe Hg(II) adsorption by kaolinite in the 0.1 ionic strength systems. The TLM is a surface coordination model developed by Davis and Leckie (1978, 1980) and modified by Hayes and Leckie (1987) to model both inner- and outer-sphere surface complex formation of cationic, anionic, and neutral solute species. This model treats the solid-solution interface as composed of two layers of constant capacitance, o-plane and ß-plane, enveloped by a diffuse layer (d-plane). Protons and other species that form inner-sphere surface complexes are adsorbed on the o-plane. Outer-sphere surface complexation occurs in the ß-plane. The diffuse layer contains loosely bound, hydrated counterions that counterbalance the charge developed in the o- and ß-planes.

Following Zachara et al. (1988), the ideal structure of kaolinite was assumed, with the ionization of {equiv}SiOH0 and {equiv}AlOH0 sites controlling surface charge. The total site density (ns) was assumed to arise from equal contributions of {equiv}SiOH and {equiv}AlOH sites (i.e., ). The kaolinite surface parameters used to predict Hg(II) adsorption were obtained from the literature and are listed in Table 1 . These literature values are specific to KGa-1 and thus pertinent to the current evaluation; although it has been shown that input parameters like specific surface, site density, and capacitance do not significantly influence the model predictions unless they are altered by an order of magnitude (Hayes et al., 1990). Surface hydrolysis constants and log Kint values of Hg(II) and other metal and ligand surface complexation reactions used in this study are listed in Table 2 . According to Sarkar et al. (1999), the Hg(II) surface complexes, {equiv}XO--HgOH+ and {equiv}XOHg(OH)2- (where {equiv}XO--M denotes outer-sphere and {equiv}XOM denotes inner-sphere surface complexation), are predicted to form in all systems and in association with silanol or aluminol sites (X is Si or Al). The adsorption constants were determined by examining the adsorption of Hg(II) by quartz ({equiv}SiOH0) and gibbsite ({equiv}AlOH0). Further, the {equiv}XOHgCl0 or {equiv}XOHgOHCl- and {equiv}XOPO3Hg2-2 surface complexes are also predicted to form in the Cl and PO4 systems. Ion association reactions and associated equilibrium constants for the formation of aqueous species of significance are listed in Table 3 . The findings of Sarkar et al. (1999) also support the formation of the aqueous species, Hg(OH)2SO2-4, Hg(OH)2H2PO-4, and Hg(OH)2HPO2-4, in addition to the well-established hydroxide, chloride, and hydroxy-chloride complexes. Conditional intrinsic equilibrium constants for various surface complexation reactions formulated with the TLM were calculated by FITEQL 3.2 (Herbelin and Westall, 1996), a computer program that combines a nonlinear least-squares fitting routine with a chemical model describing aqueous speciation and adsorption. FITEQL 3.2 was used to obtain the equilibrium adsorption constants for a given set of reactions that yields the optimum fit to the experimental adsorption data, while also satisfying the mass and charge balance constraints for all aqueous and surface chemical equilibria considered. The predicted adsorption of Hg(II) as a function of pH was then compared with that determined experimentally. A favorable comparison of the predicted and experimental results would support the validity of the chemical model (aqueous and surface species). In some systems, predicted Hg(II) adsorption deviated from that determined experimentally. When such deviations were encountered, the adsorption data were reexamined by FITEQL 3.2 to reevaluate selected surface complexation constants.


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  		Table 1 Solid and suspension properties used in the triple-layer modeling of Hg(II) adsorption by kaolinite (Riese, 1982)

 

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  		Table 2 Surface complexation reactions and intrinsic equilibrium constants [log Kint] used in the triple-layer modeling of Hg(II) adsorption.{dagger}

 

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  		Table 3 Aqueous speciation reactions and equilibrium constants used in triple-layer modeling of Hg(II) adsorption by kaolinite (Smith and Martell, 1976; Baes and Mesmer, 1986)

 

   Results and discussion
TOP
 ABSTRACT
 INTRODUCTION
 Materials and methods
 Results and discussion
Conclusions
 REFERENCES
 
Effect of pH and Ionic Strength
The adsorption of Hg(II) on kaolinite is a function of solution pH (Fig. 1) . The influence of pH on Hg(II) adsorption has been established for silica (MacNaughton and James, 1974; Bonnissel-Gissinger et al., 1999), iron oxide (Kinniburgh and Jackson, 1978), ferric hydroxide (Lockwood and Chen, 1974), bentonite (Newton et al., 1976), goethite and a sandy loamy soil (Barrow and Cox, 1992a, Barrow and Cox, 1992b), selected soils (Yin et al., 1996), and quartz and gibbsite (Tiffreau et al., 1995; Sarkar et al., 1999). Farrah and Pickering (1978) reported very little effect of pH on the adsorption of Hg(II) by kaolinite and illite in the alkaline pH range, a finding that is consistent with ours. In the present study, the adsorption of Hg(II) by kaolinite increased with increasing pH. The Hg(II) adsorption envelope is characterized by two pH values: the pH at which maximum adsorption occurs (pHmax) and the pH at which adsorption is 50% of maximum (pH50). A pH50 of 3.4 and an adsorption maxima at a pHmax of 4.4 were observed for Hg(II) adsorption from a 0.1 M NaNO3 solution. A comparison of the aqueous speciation of Hg(II) in 0.1 M NaNO3 (Fig. 2) and Hg(II) adsorption behavior indicates that the rise in Hg(II) adsorption with increasing pH corresponds to an increase in the concentration of Hg02, and to a lesser extent, HgOH+. Indeed, the pH50 value of 3.4 is consistent with the pcKa value of 3.2 for the hydrolysis reaction: (Baes and Mesmer, 1986). This observation suggests that the Hg02 species is a principal Hg(II) adsorbate, a conclusion that is consistent with that of MacNaughton and James (1974), Thanabalasingam and Pickering (1985), Lockwood and Chen (1974), and Sarkar et al. (1999).



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  		Fig. 1 Adsorption of Hg(II) on kaolinite (3.3 g L-1) as a function of pH and ionic strength (IS), and in the presence of ligands (Cl, SO4, and PO4) and metals (Ni and Pb) (IS controlled by NaNO3)

 


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  		Fig. 2 Aqueous speciation of 0.6 µM Hg(II) in (a) 0.1 M NaNO3 background and in (b) 0.1 M NaNO3 + 0.01 M NaCl background

 
According to Fig. 2, the aqueous speciation of Hg(II) does not change above pH 5, where Hg02 is the predominant solution species. However, Hg(II) retention above the pHmax gradually decreases with increasing pH. If it is assumed that Hg02 behaves as a ligand (Thanabalasingam and Pickering, 1985), then the decrease in adsorption may be explained by the loss of exchangeable ligand (H2O or OH-) from the surface (assuming inner-sphere surface complexation). Irrespective of the system, the concentration of the protonated surface functional groups will decrease with increasing pH, with a concomitant increase in the deprotonated groups. The Hg(II) adsorption process can be illustrated by the reaction:

(1)

The increased concentration of the OH- ion in solution with increasing pH may also contribute to reduced Hg(II) retention, particularly if both ligands are competing for the same surface function groups. Indeed, Mattigod et al. (1985) concluded that the decrease in boron adsorption at high pH was due to the competition between OH- and B-4 for adsorption sites. Further, OH- concentrations exceed total Hg(II) when the solution pH increases above 7.7, supporting a competition effect.

The gradual decrease in Hg(II) adsorption above pHmax (~4.4) may also be due to a decrease in HgOH+ as solution pH increases above 4.5 (Fig. 2). Through application of the TLM, Sarkar et al. (1999) suggested that the adsorption of Hg(II) by quartz and gibbsite could be explained by the adsorption of HgOH+ to the outer-sphere plane and Hg02 to the inner-sphere plane:

(2)

(3)

They also noted that the formation of the inner-sphere surface complex, {equiv}XOHgOH0, was not supported by the experimental data in either the quartz or gibbsite systems. However, the findings of Tiffreau et al. (1995) and Bonnissel-Gissinger et al. (1999) support a Hg(II) surface complex of this nature on amorphous silica and quartz. The application of Eq. [2] and [3] to explain the adsorption of Hg(II) by kaolinite (assumes the log Kint values for Eq. [2] and [3] determined for Hg(II) adsorption by quartz and gibbsite are valid for the adsorption of Hg(II) by the silanol and aluminol functional groups of kaolinite) results in a predicted adsorption envelope that describes Hg(II) adsorption at pH values above pHmax (Fig. 3) ; however, the Hg(II) adsorption edge is predicted to occur at a pH value approximately 0.5 units below the observed adsorption edge. Sarkar et al. (1999) observed a similar discrepancy when modeling the adsorption of Hg(II) by gibbsite.



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  		Fig. 3 Triple layer modeling of Hg(II) (0.6 µM) adsorption by kaolinite (3.3 g L-1) as a function of pH from a 0.1 M NaNO3 solution (open circles represent experimental data and lines represent modeled surface complexes)

 
The relative influence of the aluminol and the silanol groups in the adsorption of Hg(II) by kaolinite is evident in Fig 3. According to the TLM, Hg(II) is primarily retained by the silanol group, with both the {equiv}SiO--HgOH+ and {equiv}SiOHg-2 species dominating at low pH and the {equiv}SiOHg-2 species dominating at high pH. Approximately 15% of the total adsorbed Hg(II) is predicted to be complexed by the aluminol group (as {equiv}AlO--HgOH+) in the pH 3 to 6 range. Ward and Bassett (1990) modeled the retention of chromate by kaolinite by assuming that adsorption occurred only at the aluminol site. A similar approach used to explain Hg(II) retention by kaolinite (i.e., retention at either the {equiv}SiOH or {equiv}AlOH site alone) would under predict Hg(II) adsorption.

Adsorption of Hg(II) by kaolinite showed very little ionic strength (IS) dependance (Fig. 1). Ionic strength had no effect on the Hg(II) adsorption in the pH 4 to 7 range (encompasses the pHmax). At pH values less than 4, the Hg(II) pH50 in the 0.01 IS system shifted approximately 0.5 pH units to a lower value, relative to the pH50 in the 0.1 and 0.5 IS systems. At pH > 7, Hg(II) adsorption decreased more rapidly with increasing pH in the 0.5 IS system, relative to the 0.01 and 0.1 IS systems. Similar observations were made for Hg(II) adsorption on silica and gibbsite surfaces (Sarkar et al., 1999), although a shift in the pH50 was only observed in the gibbsite system. The adsorption data tend to support an inner-sphere Hg(II) adsorption mechanism, although such an interpretation may not be supported by the lack of response to IS adjustment (Manceau and Charlet, 1994). Further, it is evident from the application of the TLM that Hg(II) adsorption is characterized by the predicted formation of both inner- and outer-sphere complexes (Eq. [2] and [3] and Fig. 3).

Effect of Inorganic Ligands
The Hg(II) adsorption edge shifted from a pH50 of 3.4 in the absence of Cl to a pH50 of 7 in the presence of 0.01 M Cl (Fig. 1). A comparison of the adsorption pattern with the aqueous speciation of Hg(II) in the presence of 0.01 M Cl (Fig. 2) reveals that the shift in pH50 can, in part, be explained in terms of solution chemistry. According to Fig. 2, the Hg(II)-Cl complexes, HgCl02 and HgCl-3, are the dominate Hg(II) species in solution up to a pH of 7.5. At pH 6 and greater, the HgClOH0 species forms, and the Hg02 species dominates at pH values greater than 7.5. The Hg(II) adsorption and aqueous speciation data support the conclusions of others (Yin et al., 1996; Sarkar et al., 1999) that Cl reduces Hg(II) retention through the formation of soluble Hg(II)-Cl species that have little to no affinity to the variable-charge surfaces.

Mercury(II) adsorption by kaolinite, in the presence of 0.01 M Cl, was similar to that observed by quartz and gibbsite (Tiffreau et al., 1995; Sarkar et al., 1999). In addition to the predicted formation of {equiv}XO-HgOH+ and {equiv}XOHg-2 species, the TLM findings of Sarkar et al. (1999) suggest the formation of the inner-sphere surface complex, {equiv}XOHgOHCl-, according to the reaction:

(4)

However, the findings of Tiffreau et al. (1995) and Bonnissel-Gissinger et al. (1999) support the formation of XOHgCl0:

(5)

The application of Eq. [2], [3], and [4] to explain the adsorption of Hg(II) by kaolinite resulted in a predicted pH50 value that is 0.7 pH units below the experimental pH50 (Fig. 4a) . Reoptimization of the chemical model, with the kaolinite adsorption data to adjust the surface complexation constant for the formation of {equiv}SiOHgOHCl- (Eq. [4]), resulted in the predicted Hg(II) adsorption shown in Fig. 4b. A decrease in the log Kint for Eq. [4], from 2.14 to 1.53, was required to predict Hg(II) adsorption by kaolinite in the presence of Cl. The application of Eq. [2], [3], and [5] resulted in a predicted Hg(II) adsorption envelope that closely followed the experimental adsorption data, without the need to modify a surface complexation constant (Fig. 4c). Irrespective of the reaction or constant used to describe Eq. [4] for the silanol group, approximately 100% of the surface bound Hg(II) was predicted to be associated with the silanol group (principally as {equiv}XOHgCl0 or {equiv}SiOHgOHCl- and {equiv}SiOHg-2).



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  		Fig. 4 Triple layer modeling of Hg(II) (0.6 µM) adsorption by kaolinite (3.3 g L-1) as a function of pH and solution chemistry: (a) 0.1 M NaNO3 + 0.01 M NaCl with {equiv}SiOHgOHCl- formation, (b) 0.1 M NaNO3 + 0.01 M NaCl with {equiv}SiOHgOHCl- formation and log , and (c) 0.1 M NaNO3 + 0.01 M NaCl with {equiv}SiOHgCl0 formation (open circles represent experimental data and lines represent modeled surface complexes)

 
The reduced adsorption of Hg(II) in the presence of 0.01 M SO4 by Mn-oxide, quartz, and gibbsite, relative to SO4-free systems has been observed (Thanabalasingam and Pickering, 1985; Sarkar et al., 1999). However, the adsorption of Hg(II) by kaolinite in the presence of SO4 was similar to that in the absence of SO4 (Fig. 1). Since SO4 has been shown to form principally outer-sphere surface complexes with kaolinite (He et al., 1997), little impact on Hg(II) adsorption is expected. Indeed, the reduction in adsorbed Hg(II) in the presence of SO4 by Mn-oxides, quartz, and gibbsite was postulated to result from the formation of aqueous Hg(II)-SO4 complexes, specifically Hg(OH)2SO2-4. The application of Eq. [2] and [3], and the Hg(OH)2SO2-4 association reaction to explain the adsorption of Hg(II) by kaolinite results in a predicted adsorption envelope that very closely follows the experimental adsorption envelope (Fig. 5a) . According to the model, Hg(II) in the presence of SO4 is primarily retained by the silanol group, with both the {equiv}SiO--HgOH+ and {equiv}SiOHg-2 species dominating at low pH and the {equiv}SiOHg-2 species dominating at high pH. Approximately 15% of the total adsorbed Hg(II) is complexed by the aluminol group (as {equiv}AlO--HgOH+) in the pH 3 to 6 range. This distribution of adsorbed species is similar to that observed for adsorption in the absence of SO4 (Fig. 3).



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  		Fig. 5 Triple layer modeling of Hg(II) (0.6 µM) adsorption by kaolinite (3.3 g L-1) as a function of pH and solution chemistry: (a) 0.1 M NaNO3 + 0.01 M Na2SO4 and (b) 0.1 M NaNO3 + 0.01 M Na3PO4 (open circles represent experimental data and lines represent modeled surface complexes)

 
The adsorption envelope of Hg(II) in the presence of 0.01 M PO4 displayed similar pH50 and pHmax values, to those observed in the absence of PO4 (Fig. 1); however, the extent of Hg(II) adsorption at pH values equal to and greater than the pHmax was reduced in the presence of PO4. Sarkar et al. (1999) noted that PO4 decreased the retention of Hg(II) by quartz, but increased Hg(II) retention by gibbsite. For both the quartz and gibbsite systems, Hg(II) adsorption could be modeled by considering the formation of Hg(OH)2H2PO-4 and Hg(OH)2HPO2-4 aqueous complexes and the bridging surface complex, {equiv}XOPO3OHg2-2. The application of the same chemical model to explain the adsorption of Hg(II) by kaolinite results in a predicted adsorption envelope that very closely follows the experimental adsorption envelope (Fig. 5b). Again, the model indicates that Hg(II) is primarily adsorbed by the silanol group, forming the {equiv}SiO--HgOH+ and {equiv}SiOHg-2 surface species. Combined, the {equiv}SiOPO3OHg2-2 and {equiv}AlO--HgOH+ complexes account for less than approximately 15% of the total adsorbed Hg(II).

Effect of Metal Cations
It has been demonstrated that the competitive adsorption of one metal in the presence of another can be significant, altering the adsorption character of the metal (Benjamin and Leckie, 1981; Farley et al., 1985; Sarkar et al., 1999). The presence of Ni and Pb (48 and 14 µM, respectively) exerted a pronounced effect on the adsorption of Hg(II) by kaolinite (Fig. 1). The pH50 for Hg(II) adsorption by kaolinite shifted from a value 3.4 in the absence of Ni to a value 4.1 in the presence of Ni. Adsorbed Hg(II) at the adsorption maxima was also decreased by 10%, a difference that increased to approximately 30% with increasing pH above the pHmax. The inclusion of Pb similarly influenced Hg(II) adsorption. However, the pH50 was not altered by the presence of Pb.

The chemical model employed to predict Hg(II) adsorption by kaolinite in the presence of Ni and Pb included Eq. [2] and [3], and accounted for the formation of {equiv}XONiOH0 and {equiv}XOPbOH0 for both the silanol and aluminol groups (after Sarkar et al., 1999). The predicted Hg(II) adsorption envelope, in the presence of Ni, successfully modeled the experimental data at pH values above pHmax (Fig 6a) . Similar to the model prediction for Hg(II) adsorption in the absence of metals and ligands, the adsorption edge was predicted to occur at a lower pH value than experimentally observed. The predicted Hg(II) adsorption envelope in the presence of Pb was significantly different than the experimental envelope (Fig. 6b). In both the Ni and Pb systems, the formation of the {equiv}SiO--HgOH+, {equiv}AlO--HgOH+, and {equiv}SiOHg-2 species was predicted to account for approximately 100% of the adsorbed Hg(II); however, as discussed by Sarkar et al. (1999), the inability of the TLM to predict adequately the pH50 and pHmax of the Hg(II) adsorption envelope in the presence of Ni or Pb indicates that the adsorption behavior of metals by oxide surfaces is more complex than can be addressed by the TLM.



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  		Fig. 6 Triple layer modeling of Hg(II) (0.6 µM) adsorption by kaolinite (3.3 g L-1) as a function of pH and solution chemistry: (a) 0.1 M NaNO3 + 48 µM Ni(NO3)2, (b) 0.1 M NaNO3 + 14 µM Pb(NO3)2 (open circles represent experimental data and lines represent modeled surface complexes)

 

   Conclusions
TOP
 ABSTRACT
 INTRODUCTION
 Materials and methods
 Results and discussion
 Conclusions
REFERENCES
 
The adsorption of Hg(II) by kaolinite was primarily influenced by pH and the presence of Cl, Ni, and Pb. Ionic strength and the presence of SO4 and PO4 had relatively little impact on the Hg(II) adsorption envelope. The general form of the Hg(II) adsorption envelope was predicted through the application of the TLM by assuming that the kaolinite surface functional groups were equally distributed between the silanol and aluminol groups. It was further assumed that all equilibrium constants used to describe the retention of Hg(II), as well as the other metals and ligands examined, by quartz (silanol) and gibbsite (aluminol) surfaces were applicable to the silanol and aluminol sites on the kaolinite surface. In general, the Hg(II) adsorption envelopes could be described without modification of the model parameters. Prediction errors were primarily associated with the pH50 values, although adsorbed Hg(II) at the pHmax in the presence of Pb was severely under-predicted. The model also suggests that the silanol group was responsible for retaining the bulk of the adsorbed Hg(II) (>85%), through the formation of the {equiv}SiO--HgOH+ and {equiv}SiOHg-2 species (and the {equiv}XOHgCl0 or {equiv}XOHgOHCl- species in the Cl system). Only in the Ni and Pb systems was the {equiv}AlO--HgOH+ complex predicted to play an important role [25–35% of the total Hg(II) retained]. Despite the relatively small impact of the aluminol group on Hg(II) retention, it is clear that both the {equiv}SiOH and {equiv}AlOH sites must be considered when addressing the adsorption of Hg(II) by kaolinite.


   ACKNOWLEDGMENTS
 
The authors express their appreciation to Mr. David Bass of the Perkin-Elmer Corporation for the use of a flow injection mercury spectrophotometer. Financial and instrumental support from the Geological Society of America, and The University of Tennessee, Department of Geological Sciences, Institute of Agriculture, and the Office of Academic Research is gratefully acknowledged.

Received for publication October 22, 1999.


   REFERENCES
TOP
 ABSTRACT
 INTRODUCTION
 Materials and methods
 Results and discussion
 Conclusions
 REFERENCES
 




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