Environ. Sci. Technol. 2001, 35, 1953-1958
Determination of Silver Speciation in Natural Waters. 1. Laboratory Tests of Chelex-100 Chelating Resin as a Competing Ligand
competition between the added ligand (LA) and the natural ligands (L) present in the sample for the metal of interest (M). This is often described as a competing ligand exchange approach and is described by the equilibrium expressions in eqs 1 and 2:
wMa+ + xLb- S Mw(L)x(a-b)
(1)
RUSSELL T. HERRIN, ANDERS W. ANDREN, AND DAVID E. ARMSTRONG* Water Chemistry Program, University of WisconsinsMadison, 660 North Park Street, Madison, Wisconsin 53706
yMa+ + zLAc- S My(LA)z(a-c)
(2)
Batch and column experiments were performed to investigate the suitability and chemical characteristics of Chelex-100 for use as a competing ligand in ionic silver (Ag(I)) speciation determinations in natural waters. A conditional stability constant (Kcond) for Ag+ chelation by iminodiacetate groups on the surface of Chelex resin was determined by fitting results of batch and column experiments with an equilibrium speciation model. Results of experiments in which Chelex competed with cyanide ion and thiosulfate ion for aqueous Ag+ were fitted well by a model in which log Kcond(Ag-Chelex) was set to 7.2. This value is similar to literature equilibrium constants for a 1:1 Ag+-EDTA chelate. In batch experiments with Chelex, equilibration times of 24 h were found to be sufficient to bring samples close to equilibrium. Effects of resin counterion and total Ag(I) concentration on extent of Ag(I) chelation were found to be minor. Effect of pH on Ag(I) chelation was minor over a range of 6-10. Column experiments (detention time ) 6 s, empty-column basis) in which thiosulfate competed with Chelex for Ag(I) gave similar results to batch experiments with thiosulfate. This implies that batch and column experiments could be compared to explore ligands in natural water systems with different rates of dissociation.
Introduction Silver, particularly in its ionic form (Ag(I)), is of environmental interest because of its demonstrated toxicity to aquatic organisms (e.g., refs 1 and 2). As has been hypothesized for other metals, it is likely that the toxic effects of Ag(I) may change if it is present primarily as a dissolved complex rather than a free (hydrated) cation (3). Fresh and brackish surface waters contain a broad range of organic and inorganic species with a range of complexing strengths for silver and other metals. These ligands may be in the form of simple organic or inorganic compounds, or they may be functional groups on large, complex organic molecules or clusters of molecules. The large number of chemically distinct ligands, particularly organic ligands, in surface waters has led to the development of indirect methods of determining the effective strength and concentration of the combination of ligands found in a given water. Often these methods are based upon addition of a ligand of known binding strength and observation of the * Corresponding author e-mail:
[email protected]; phone: (608)262-0768; fax: (608)262-0454. 10.1021/es001509x CCC: $20.00 Published on Web 04/06/2001
2001 American Chemical Society
Differences between the observed concentration of My(LA)z and the concentration predicted by equilibrium speciation modeling, based on a known formation constant (KMLA), provide information about the effective strength (KML) and concentration of the Mw(L)x complex. Most frequently, KML and concentration of L are quantified by titrating a sample (to which LA has been added) with the metal of interest. The titration curve is linearized using an algebraic rearrangement of the Langmuir isotherm. Estimates of KML and the concentration of L are calculated from the slope and intercept of this line (4-6). Speciation modeling can however be very complicated; many complexes, both dissolved and colloidal, and many complex stoichiometries may be present in a system simultaneously. In addition, other metals present in a sample may compete with the metal of interest for the added ligand. Competing ligand experiments require an analytical method for determination of the concentration of the My(LA)z complex or complexes. For metals other than silver, a common strategy has been to choose a competing ligand that forms a surface-active complex with the metal of interest. The complex My(LA)z can then be sorbed to a mercury drop for electrochemical determination of the complexed metal (e.g., adsorptive cathodic stripping voltammetry, ACSV) (4, 5). In another method, the complex is extracted into a nonpolar organic phase such as chloroform (6). Problems exist in both of these approaches to silver speciation experiments in freshwaters. The most effective electrodes for use in ACSV experiments are mercury drop electrodes. Unfortunately, the reduction potentials of Ag+ (E° ) 0.7996 V relative to SHE), Hg22+ (E° ) 0.7973 relative to SHE), and Hg2+ (E° ) 0.851 V relative to SHE) are very close to each other (7). The high mercury reduction currents that result from using a mercury-based electrode mask signals from Ag(I) reduction. Other electrode materials can be used for detection of silver, but detection limits are high relative to environmental levels (8), and electrode fouling and other matrix effects are problematic (9, 10). As a result, ACSV appears to be a poor choice for detection of silver complexes. Liquid/liquid extraction methods for separating My(LA)z have been used effectively to investigate silver speciation in ocean and estuarine waters (6) and in wastewaters and their (fresh) receiving waters (11). The ammonium salt of diethyldithiocarbamate (DDC) was used as a competing ligand in these investigations. The 1:1 complex of Ag+ and DDC has a stability constant of 109.1 (11). This silver titration/extraction procedure can quantify silver-natural ligand stability constants when RAgDDC (RAgDDC ) [AgDDC]/[Ag+]) is within an order of magnitude of RAgL (6). The utility of this technique may be limited in freshwater systems with high dissolved organic matter concentrations since some silver may be extracted into the organic phase in the form of extractable Ag-natural ligand complexes. When this occurs, it is impossible to distinguish the fraction of Ag(I) extracted with the competing ligands from the fraction extracted with natural ligands. The VOL. 35, NO. 10, 2001 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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extent of this artifact may vary from system to system, but it is likely to be a problem in rivers or lakes high in dissolved and colloidal organic matter. The functional group on a chelating or ion-exchange resin represents another potentially useful competing ligand. In this case, separation of My(LA)z is accomplished by separating the resin from the solution by filtration and determining the concentration of metal eluted from the resin or remaining in solution. Resins (most often Chelex-100, Bio-Rad, Inc.) have been used in speciation of Cu, Ni, Cd, Pb, and Zn (4, 12, 13). Measurements were performed by passing solutions through columns of Chelex resin at moderate to fast flow rates. Under these conditions, complexes that are labile within column residence times will dissociate, and the metal ion will bind to the resin. Complexes that are not labile under these conditions of resin biding strength and column detention time will dissociate. For example, colloidal complexes of a diameter greater than the resin pore diameter will likely to exit the column intact. Previous experiments (4, 12, 13) were performed in conjunction with other speciation techniques to determine separate concentrations of ligands that strongly bind metals and of ligands that are not as strong but do not dissociate during the Chelex-solution contact time. Chelex-100 resin may also be useful in experiments in which resin-sample contact times are relatively long, and the effects of dissociation kinetics of solution complexes on fractionation of metal between ligands in solution and the resin are minimized (i.e., the resin-solution system is allowed to approach an equilibrium condition). Experiments of this type have been performed before (14, 15), but only to determine the concentration of “stable” or “inert” complexes of the metal of interest. The binding strength of the resin functional groups were not considered, so an analysis like that represented by eqs 1 and 2 was not performed. In none of these previous investigations was the complexation of Ag(I) studied. Chelex resin has several useful characteristics with respect to Ag-competing ligand equilibration experiments. Relative to many other “heavy metal” ions (e.g., Pb2+, Hg2+, Cu2+, and Zn2+), Chelex has a relatively low affinity for Ag+. As a result, the resin can be added in relatively large, easily measurable quantities and still compete with low concentrations of ligands with stronger silver-binding properties. This large effective concentration of Chelex functional groups also leads to a small fraction of groups occupied by strongly bound trace metals such as those mentioned above, and corrections do not need to be made for competition from these metals. At the same time, Chelex resin has a much stronger affinity for Ag+ than for major cations such as Na+, Mg2+, and Ca2+, so competition from these ions is likely to be negligible. Finally, use of Chelex resin in competing ligand experiments involves fewer experimental difficulties and hazardous reagents than liquid/liquid extraction, and Chelex is therefore more adaptable to performance on large numbers of samples and to use in the field. Results of Chelex equilibrations can be used in conjunction with results of short-term exposure (column) experiments and possibly with results of liquid/liquid extraction experiments to quantify the effects of natural ligands over a broad range of binding strengths. Iminodiacetate groups bound to a solid resin matrix cannot be assumed to have the same metal sorption characteristics as, for example, dissolved alkyliminodiacetates, and thus the properties of the resinbound groups must be characterized before the results of field experiments can be interpreted. In this paper, we investigate the effects of Chelex-solution contact time, pH, resin counterion, and dissolved ligands on Ag-Chelex interactions in batch systems. We use the results to determine the effective Ag(I)-Chelex complexation constant under a variety of conditions. Our goal is to characterize Ag(I)-Chelex interactions to an extent sufficient to allow Chelex-100 to be 1954
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used effectively in field investigations including competing ligand equilibration experiments.
Experimental Section Chelex-100 Resin: Properties and Preparation. Chelex-100 (referred to simply as Chelex hereafter) is composed of 100mesh beads of a styrene-divinylbenzene copolymer. Methyliminodiacetate (IDA) ions (CH3-N-(CH2COO-)2) are bound to the polymer surface. These groups coordinate metals through both oxygen and nitrogen bonds, similar to ethylenediaminetetraacetic acid (EDTA). This structure makes Chelex much more selective for heavy metals and alkaline earth metals than for ions of alkali metals such as Na+ and K+ (16). In particular, the ability of the nitrogen atom of the IDA group to share lone pair electrons with metals leads to much stronger metal-IDA complexes than, for example, complexes of alkylmalonic or alkylacetic acids with ions of metals such as Cu, Ni, Zn, Co, Cd, and Ag (17). The concentration of IDA groups is 0.4 mmol (mL of resin)-1 when the resin is hydrated and in the sodium form (Bio-Rad, Inc.), or 1.6 mmol (g of resin)-1 when the resin is dry and in the hydrogen form (18). Protonation constants have been experimentally determined for the resin; the values of pKA1 and pKA2 have been reported as 3.2 and 9.12, respectively (18, 19). The styrene-divinylbenzene copolymer of which Chelex resin is composed has a cross-linked, highly porous structure. The average size of these pores is approximately 1.5 nm (although this pore size varies with pH). Most of the surface area of Chelex resin is in this pore space rather than on the surface of the resin, and as a result, the outer surface of the resin has been found to contribute little to the chelating capacity of Chelex (16). Thus, the IDA groups on Chelex will not compete directly for metals bound to colloids too large to enter the pore structure. It is likely, however, that in batch experiments the resin will compete with colloids by reducing concentrations of ionic or hydrated Ag+ and thereby driving Ag(I) dissociation from colloids to maintain equilibrium. While the behavior of Chelex resin with metals in natural water samples has been studied extensively, only a few of these studies have addressed Ag-Chelex interactions (20, 21). Equilibrium constants for dissolved organics containing an IDA functionality have however been determined. Values of log β1 and log β2 for formation of Ag-L and Ag-L2, where L is n-propyliminodiacetic acid, are 4.29 and 7.5, respectively; for n-butyliminodiacetic acid, these values are 4.36 and 7.6, respectively (17). Stability constants for Ag+ association with IDA groups bound to Chelex have not been determined, but based on work with Ca2+, Cu2+, and other metals (18), they will likely be similar to those of dissolved alkyliminodiacetates at pH levels sufficiently high that the IDA groups are at most monoprotonated. This will occur at pH values higher than 3.3, the pKA1 of the resin (22). Chelex resin was purchased from Bio-Rad, Inc. in the Na+ form. For our experiments, Chelex was used with either Na+ or Ca2+ as the counterion on the IDA groups. A total of 10-20 g of resin was weighed into a Teflon bottle, and trace metal contamination was removed from the resin by mixing it for 30 min with 125 mL of 2.5 M HNO3 (Ultrex grade; J. T. Baker, Inc.). The resin and acid were mixed on an incubator-shaker set at 100 rpm and 25 °C. The resin was separated from the acid solution by filtration and converted to the Na+ form by mixing for 2 h with 125 mL of 2 M NaOH (SigmaUltra grade, Sigma Chemical). For use in the Na+ form, the resin was separated from the NaOH and mixed for 2 h with 125 mL of a NaC2H3O2-HC2H3O2 buffer (pH 5.6) (23). To obtain the Ca2+ form, the resin was mixed for 2 h with 125 mL of 2 M Ca(NO3)2 instead of the buffer solution. In both cases, as much solution as possible was removed from the resin by vacuum filtration, and aliquots were weighed into Teflon
vials, capped tightly, and stored at room temperature. To determine the moles of functional groups in a given mass of the hydrated resin, experiments were performed in which measured masses of the Na+ and Ca2+ form of Chelex were converted to the H+ form, dried, and weighed. These experiments indicated that a given quantity of resin in the Na+ form weighs two times as much as the same quantity in the Ca2+ form and occupies approximately two times the volume. Sodium and calcium ions form complexes, not chelates, with the carboxylic acid groups in Chelex-bound iminodiacetic acid. At high concentrations of Na+, Na2IDA predominates, and at high Ca2+ concentrations, CaIDA predominates. Thus, the mass of cations in the sodium form of the resin is approximately double the mass of the calcium form, as is the volume occupied by cations. To prevent contamination, Chelex preparation was performed in a clean room environment. Mixing steps were not performed in the clean room, but during these steps, reagents were sealed in bottles, and the bottles were sealed in three new polyethylethylene bags. Bottles, vials, filters, and filtration apparatuses were constructed of Teflon and cleaned prior to use with a multistep acid cleaning procedure similar to that of Shafer et al. (24). High-purity water (Milli-Q, Millipore, Inc.; or Nanopure, Barnstead, Inc.) was used in all rinsing steps and in preparation of solutions. Competition Experiments. Solutions used in Chelex competition experiments were made up in 500-mL Teflon or glass bottles. Despite the buffering step performed during preparation of the Na+ form of Chelex, the pH of poorly buffered water to which this form was added increased to values greater than 10. The high pH was likely the result of replacement of Na+ by H+ on the resin functional groups, which has also been described as the hydrolysis of the Na+ form of Chelex (25). Chelex resin has a much higher affinity for H+ than for Na+, as evidenced by the need to use NaOH as opposed to, for example, NaCl to convert the resin from the hydrogen to the sodium form. Our goal was to perform experiments at pH values approximating those in wellbuffered surface freshwaters, so 60 g of the acetate buffer used in resin preparation was included as part of the 500-mL test solutions to which the Na+ form of the resin was to be added ([C2H3O2]T ) 12 mM). Calcium ions have a much higher affinity for Chelex than Na+, as evidenced by the fact that a solution of Ca(NO3)2 in contact with the sodium form of the resin will convert it to the calcium form. Solution pH did not increase appreciably when the Ca2+ form of the resin was added to test solutions with a low buffering capacity. Solutions in which the Ca2+ form of the resin was used were buffered with 1 mM NaHCO3. Experiments indicated that the choice of Na+ or Ca2+ as a counterion had negligible effect on the extent of Ag+ chelation by Chelex. Acetate and bicarbonate buffers were chosen becausesunlike buffers such as Tris and maleic acidsacetate, bicarbonate, and carbonate ions have very low stability constants for complexation with Ag+ (26). Emphasis was placed on studying systems in which sample solutions and Chelex were equilibrated for long periods (batch experiments). Known concentrations of dissolved ligands, silver standard (dissolved AgNO3; High Purity Standards, Inc.), and Chelex were added to the acetate- or bicarbonatebuffered solutions. Titrations with silver or dissolved ligands were performed by preparing a separate bottle for each titrant concentration. Cyanide and thiosulfate were the principal ligands used to compete with Chelex for Ag+. Silver was added at concentrations ranging from 0.47 to 9.28 nM, and the mass of Chelex added led to an effective concentration of IDA groups ranging from 0.3 to 1 mM. Samples were capped, triple-bagged, and mixed for 23-25 h on an incubator/shaker (model 236, Fisher Scientific) set at 25 °C and 100 rpm. Resin was separated from the solution phase using all-Teflon or
all-glass resin columns. Solution-phase samples were collected for pH determination and for Ag concentration analysis. To the latter, sufficient concentrated HNO3 (Ultrex grade, J. T. Baker, Inc.) was added to make the concentration 0.15 M (1% HNO3, v/v). Silver chelated with the resin was eluted using 15 mL of 2 M HNO3 followed by 15 mL of ultrapure water. Second elutions and mass balances indicated that this technique quantitatively eluted Ag from the resin. To determine whether Chelex kinetics were fast in the presence of free Ag+ and fast-dissociating dissolved Ag ligands, experiments were also performed in which solutions were contacted with a column of Chelex for relatively short time periods (column experiments). Solutions were pumped through columns of Chelex at 6-7 mL min-1, leading to a detention time of approximately 6 s (empty-column basis). Resin columns used for column tests are approximately 12 mm in diameter, and 0.25 g of resin did not dependably cover the entire surface of the column frit. To avoid incomplete solution-Chelex contact, 0.5 g of resin was used in column experiments. Columns of 0.5 g of resin were 5-6 mm in length, and in all cases completely covered the resin column frit. Silver was eluted from the column, and samples were taken for solution Ag and pH determination by the same methods as in batch experiments. Sample preparation and elution steps were performed using equipment cleaned as described in the resin preparation section. Preparation and elution steps were performed in a clean room environment. No pattern of Ag contamination was observed, indicating that these precautions were sufficient. Some experiments were performed using Teflon bottles and resin columns, and others were performed in glass bottles and columns. In both materials, contamination was not detected, and Ag recovery was acceptable (27). Experiments involving Ag(I) require special care because the ion may be reduced to Ag(0) by exposure to visible light. Precautions were therefore taken to prevent photoreduction of Ag(I). Samples were kept in opaque bags during equilibration steps, and resin separations were performed under the darkest conditions practical. Analysis of Ag Concentration. Silver concentrations in eluates and solution-phase samples were analyzed using a graphite furnace atomic absorption spectrophotometer (GFAAS) equipped with Zeeman background correction and pyrolytically coated graphite tubes containing L’vov platforms. Ammonium dihydrogen phosphate was used as a matrix modifier. The instrument is situated in a trace metal clean laboratory. Sensitivity of this method was improved by pipetting and evaporating 2-3 35-µL vol of sample onto the platform, thus increasing the mass of Ag atomized. Quality assurance procedures included frequent analysis of standard and blank solutions and frequent matrix spikes. Typical matrix spike recoveries ranged from 85 to 110%. In some experiments, the concentration of Ag in the solution phase was too low to be analyzed directly by GFAAS. In these cases, Ag in samples was preconcentrated by coprecipitation with cobalt pyrrolidinedithiocarbamate (28, 29). Speciation Modeling. Modeling of equilibrium speciation in solution was performed using MINEQL+ version 4.0 (Environmental Research Software, Inc.). Stability constants for dissolved species were taken from the MINEQL+ library or from Martell and Smith (26). Stability constants for resinbound species were taken from Pasavento and Biesuz (22). Important stability constants used in modeling are shown in Table 1. Chelex resin was modeled as if its surface functional group was a dissolved species. A concentration of 1.6 mmol of IDA/g of dry resin in the hydrogen form was used in determining effective IDA concentrations in solutions containing Chelex (18). VOL. 35, NO. 10, 2001 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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TABLE 1. Stability Constants of Some Important Complexes for Modeling Competitive Behavior of Chelex and Dissolved Ligandsa ligand (L)
cation (M)
CN-
Ag+
CNCN-, OHCNS2O32S2O32S2O32S2O32S2O32S2O32S2O32IDAChelex IDAChelex IDAChelex IDAChelex methyl-IDAdissolved propyl-IDAdissolved propyl-IDAdissolved EDTA
Ag+ Ag+ H+ Ag+ Ag+ Ag+ H+ H+ Ca2+ Na+ H+ H+ Ca2+ Ca2+, H+ Ca2+ Ag+ Ag+ Ag+
equation
log K
K ) [ML2 K ) [ML3]/[M][L]3 K ) [M(OH)L]/[M][OH][L] K ) [ML]/[M][L] K ) [ML]/[M][L] K ) [ML]/[M][L]2 K ) [ML]/[M][L]3 K ) [ML]/[M][L] K ) [M2L]/[M]2[L] K ) [ML]/[M][L] K ) [ML]/[M][L] K ) [ML]/[M][L] K ) [M2L]/[M]2[L] K ) [ML]/[M][L] K ) [M(HL)2]/[M][H]2[L]2 K ) [ML2]/[M][L]2 K ) [ML]/[M][L] K ) [ML2]/[M][L]2 K ) [ML]/[M][L]
20.48a 21.70a 13.22a 8.88a 8.82a 13.50a 14.20a 1.60a 0.60a 1.90a 0.60a 9.12b 12.42b 4.40b 19.76b 7.36a 4.29c 7.5c 7.2a
/[M][L]2
a Constants are taken or calculated from Martell and Smith (26). Constants are taken or calculated from Pasavento and Biesuz (22). c Constants are taken or calculated from Israeli and Pettit (17). b
FIGURE 1. Time course of Ag+ uptake by Chelex, measured as loss from solution. Chelex was used in the Na+ form (1.5 g/sample).
Results and Discussion Kinetics of Ag Uptake by Chelex. Iminodiacetate groups on the surface of Chelex resin are minor in number as compared to those within the pores of the resin beads (16). As compared to dissolved iminodiacetate, the localization of functional groups and the tortuous paths represented by the resin pores are likely to greatly increase the importance of mass transfer as a limiting step in Ag(I) chelation kinetics. While quantification of the kinetics of Ag-Chelex interactions was not a primary goal of this research, contact times with Chelex sufficient to produce near-equilibrium conditions were needed. An experiment was performed in which batch samples were set up with equal concentrations of Ag and Chelex (9.2 nM and 1.5 g of Na+ form, respectively) and increasing concentrations of CN-. Samples of the aqueous phase for Ag concentration determination were taken 2, 7, 25, and 29 h after Chelex addition (Figure 1). Results of tests at all CN- concentrations indicate that solution-phase Ag concentrations are reasonably stable after 25 h. The sample that contained 20 nM CN- appears to still be approaching equilibrium between 25 and 29 h. If equilibrium was not reached, our estimate of the conditional stability constant for chelation of Ag(I) by Chelex will be slightly lower than the true thermodynamic constant. The rate of change is very slow however, and longer equilibration times might also lead 1956
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FIGURE 2. Effects of initial solution concentration of Ca2+ on pH and resin uptake of Ag+. Experiments were 24-h batch studies. Chelex was used in the Na+ form (0.5 g/sample). No buffer was added. to increased oxidation or hydrolysis of ligands in natural waters. Thus, an equilibration time of 23-25 h was considered acceptable for these experiments; as shown later, 23-25-h equilibration times gave reproducible results. Effects of pH and Resin Counterion. On the basis of measured pKA values for Chelex resin (18, 19), hydrogen ion would be expected to compete with Na+, Ca2+, and Ag+ for IDA groups on Chelex resin. As solution pH is reduced from 10 to 6, modeling predicts that available hydrogen ions would replace Na+ and Ca2+ ions bound to the resin and would also replace appreciable quantities of resin-bound Ag+. Batch experiments indicated however that the effect of pH on Ca2+ and Ag+ association with the resin was negligible. Hydrogen ion effectively displaced Na+ on resin functional groups, as evidenced by extreme pH increases when the Na+ form of the resin was introduced to poorly buffered solutions. Changes of this magnitude were not observed when the Ca2+ form of the resin was introduced to poorly buffered systems nor when the Na+ form of the resin was introduced to a solution containing millimolar concentrations of CaNO3 (Figure 2). When the Na+ form of the resin was added to an unbuffered, Ag(I)-containing solution (Figure 2, leftmost bar), the pH increased to a value over 10 during equilibration. When Ca2+ was added to the same system, however (four rightmost bars, Figure 2), pH values were approximately 6. Some drift occurred in pH values, presumably a result of low ionic strength and dissolution of atmospheric CO2 during measurement. Changes in resin volume upon introduction of CaNO3 indicate that the resin is converted to the Ca2+ form and that chelation of Ca2+ by the resin limits complexation of H+. Modeling of the Chelex-Ca2+ system using literature stability constants (22, 26) predicts a higher pH than was observed. The difference may reflect uncertainty in the stability constant for the Ca(IDA)2 complex. The stability constant was calculated for dissolved methyliminodiacetic acid, as opposed to resin-bound methyliminodiacetic acid. It is also possible that the need to maintain a charge balance in the resin pores led to a different hydrogen ion activity in the pores than that in the bulk solution. Neither the form of the resin nor the solution pH has an appreciable effect on the fraction of Ag+ complexed by the resin (Figure 2). This conclusion is supported by comparison of the results of several experiments in which a Ag-Chelex system was titrated with CN- (Figure 3). Different counterions and solution pH values in these experiments did not have strong effects on the competition between Chelex and CNfor Ag+. In addition, the results in Figure 3 indicate that, over a range from 0.46 to 4.6 nM, the total concentration of Ag(I) did not appreciably affect the fraction of Ag(I) chelated by the resin. Despite the stability of the system with respect to
FIGURE 3. Results of batch titrations of Ag-Chelex systems with CN-. When Chelex was used in the Na2+ form, 0.5 g of resin was used; when it was used in the Ca2+ form, 0.25 g of resin was used. An asterisk (*) indicates that the sample was excluded due to apparent contamination. Two asterisks (**) indicate that the sample was excluded due to apparent losses to bottle. Three asterisk (***) indicate that the 100 nM CN- titration step was not performed in this experiment. pH and Ag(I) concentration, we recommend performing experiments on field samples at Ag(I) concentrations and pH values close to those in the unaltered sample. Competition between Chelex Resin and Dissolved Ligands. The Ag+ chelation behavior of Chelex in the presence of dissolved ligands was investigated using two different dissolved ligands: cyanide ion (CN-, introduced as NaCN) and thiosulfate ion (S2O32-, introduced as Na2S2O3). These ligands were chosen for their contrasting stability constants for complexation with Ag+ and for their resistance to oxidation relative to reduced sulfur compounds (e.g., sulfide and organic thiols), another class of strong Ag+-complexing agents. In addition, stability constants for complexes of these ligands with Ag+ and other important species were available (see Table 1). No literature stability constants are available for chelation of Ag+ by resin-bound IDA, so a conditional stability constant (Kcond) was determined by adjusting this value in an equilibrium speciation model and comparing modeling results to results of competing ligand titration experiments. This constant is referred to as conditional because it cannot be demonstrated to apply at conditions other than those used in the experiments performed to determine it and is therefore not a true thermodynamic equilibrium constant. Since a variety of experiments indicated that pH has negligible effects on Ag-Chelex interactions over a broad pH range, pKA values for Chelex were omitted from the model. Several experiments were performed in which CN- was the titrant (see Figure 3). Throughout the range of CNconcentrations used in the equilibrations, the Ag(CN)2complex was present at concentrations at least 1000 times higher than any other Ag+-CN- complex (see Table 1). Mean Chelex-bound Ag fractions at a given CN- concentration were calculated from these results and compared to modeled results. A conditional Ag-Chelex stability constant of 107.2 was found to provide an acceptable fit to the CN- results (Figure 4). Because of the modeling strategy used, it is difficult to describe the quality of model fit to the data in statistical terms. Modeled results based on other stability constants are included in Figure 4 however to demonstrate that 107.2 provides a good fit to the data relative to other values. This constant is very close to the stability constant for AgL2 complexes, where L is n-propyliminodiacetic acid (log K ) 7.5) and n-butyliminodiacetic acid (log K ) 7.6) (26). It is even closer to the stability constant for a 1:1 Ag-ethylenediaminetetraacetic acid (EDTA) complex (log K ) 7.2) (26).
FIGURE 4. Comparison of CN- batch titration results (mean values from Figure 3) to a speciation model in which the equilibrium constant for Ag complexation by Chelex was set to 107.2. Results based on other equilibrium constants are included for comparison. Error bars represent 1 standard deviation of the mean.
FIGURE 5. Equilibrium speciation modeling results for competition of Chelex and S2O32- for Ag+. Conditions are similar to those in S2O32- batch titration experiments: 0.3 mM resin-bound IDA in Ca2+ form, 1 mM NaHCO3 in solution, pH 8.9. The equilibrium constant for Ag complexation by Chelex was set to 107.2. EDTA contains two iminodiacetate groups and therefore would be expected to complex Ag+ to an extent similar to paired iminodiacetate ions. These similarities indicate that a single Ag+ ion is complexed by two IDA groups when it is within the resin pore structure. An experiment was performed with thiosulfate as the titrant rather than CN-. Modeling revealed that, over the range of S2O32- concentrations studied, the predominant Ag+-S2O32- complex changed from AgS2O3- to Ag(S2O3)23(Figure 5). Despite this complexity, however, the three titration points from the S2O32- batch experiment (Figure 6) fit well with the model in which the Ag-Chelex stability constant was set at 107.2, as developed in the CN- titrations. A pair of column experiments was also performed using S2O32as the competing ligand, and the results were in agreement with the model developed in batch experiments (27). It appears that the extremely high effective concentration of Chelex in a packed column overcomes the mass-balance rate limitation observed in batch systems (see Figure 1). In a second experiment with S2O32- as the dissolved ligand, the pH was forced to 6.6 with acetate buffer and Ag+ was the titrant. These results gave a very poor model fit when the pH was forced to 6.6 (i.e., when the model predicted competition between Ag+ and H+ for resin sites) but provided a good fit when the pH was raised to 9. This reinforces the observation, noted above, that H+ has a small effect on Ag+ chelation relative to that predicted by literature constants. The conditional equilibrium constant for Ag chelation by Chelex has been defined through competition with dissolved VOL. 35, NO. 10, 2001 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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Sea Grant College Program, National Oceanic and Atmospheric Administration, U.S. Department of Commerce, and from the State of Wisconsin, Federal Grant number NA46RG0481, project number R/MW-77.
Literature Cited
FIGURE 6. Comparison of S2O32- batch titration results to the speciation model shown in Figure 5. Measured pH in the batch solutions was 8.9.
FIGURE 7. Family of curves defining natural ligand concentration and Ag binding strength pairs detectable by Chelex competing ligand experiments. Curves are shown for five different fractions of total Ag associated with Chelex. ligands such as CN- and the effects of pH and major ions determined. Thus characterized, Chelex can be used in quantifying the binding strength and concentration of Ag ligands of unknown structure in aquatic systems. For a known mass of Chelex, pairs of natural-ligand formation constants (K(AgL)) and total concentrations ([L]tot) have been predicted for different fractions of total Ag associated with Chelex (Figure 7). Calculations are based on an Ag mass balance and mass action laws for formation of Ag-Chelex and AgL. For this analysis it has been assumed that only one Ag-binding ligand is present, that the ligand forms only 1:1 complexes with Ag, and that pH and major ions do not affect Ag interactions with Chelex or the natural ligand. It is clear from the plot in Figure 7 that ligands with a wide range of Agbinding strengths and over a wide range of concentrations can be detected using Chelex as a competing ligand. It is also clear that batch Chelex experiments alone cannot define a unique binding strength-concentration pair. If this is the desired result, titration with Ag+ or comparison to results from a different competing ligand is required.
Acknowledgments This research was supported by the National Association of Photographic Manufacturers and by the University of Wisconsin Sea Grant Institute under grants from the National
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(1) Fisher, N. S.; Hook, S. E. Proceedings of the Fifth International Conference on Transport, Fate and Effects of Silver in the National Environment; Andren, A. W., Bober, T. W., Eds.; University of Wisconsin Sea Grant; 1997. (2) Webb, N. A.; Wood, C. M. Environ. Toxicol. Chem. 1998, 17 (4), 579-588. (3) Campbell, P. G. C. In Metal Speciation and Bioavailability in Aquatic Systems; Tessier, A., Turner, D. R., Eds.; John Wiley & Sons: New York, 1995. (4) Donat, J. R.; Lao, K. A.; Bruland, K. W. Anal. Chim. Acta 1994, 284, 547-571. (5) Xue, H. B.; Kistler, D.; Sigg., L. Limnol. Oceanogr. 1995, 40 (6), 1142-1152. (6) Miller, L. A.; Bruland., K. W. Anal. Chim. Acta 1994, 284, 573586. (7) CRC Handbook of Chemistry and Physics; Lide, D. R., Ed.; CRC Press: Boca Raton, FL, 1992. (8) Wang, J.; Ruiliang, L.; Huiliang., H. Electroanalysis 1989, 1, 417421. (9) Song, S.; Fedkiw, P. S. Proceedings of the Third International Conference on Transport, Fate and Effects of Silver in the National Environment; Andren, A. W., Bober, T. W., Eds.; University of Wisconsin Sea Grant; 1995. (10) Dao, P. T.; Robillard, K. A.; Schildkraut, D. E. Proceedings of the Third International Conference on Transport, Fate and Effects of Silver in the National Environment; Andren, A. W., Bober, T. W., Eds.; University of Wisconsin Sea Grant; 1995. (11) Adams, N. W. H.; Kramer, J. R. Environ. Toxicol Chem. 1999, 18 (12), 2674-2680. (12) Sedlak, D. L.; Phinney, J. T.; Bedsworth, W. W. Environ. Sci. Technol. 1997, 31, 3010-3016. (13) Figura, P.; McDuffie, B. Anal. Chem. 1979, 51 (1), 120-125. (14) Hart, B. T.; Davies, S. H. R. Aust. J. Mar. Freshwater Res. 1977, 28, 397-402. (15) Figura, P.; McDuffie, B. Anal. Chem. 1980, 52, 1433-1439. (16) Apte, S. C.; Batley, G. E. In Metal Speciation and Bioavailability in Aquatic Systems; Tessier, A., Turner, D. R., Eds.; John Wiley & Sons: New York, 1995. (17) Israeli, M.; Pettit, L. D. J. Chem Soc. Dalton Trans. 1975, 414417. (18) Pasavento, M.; Biesuz, R.; Gallorini, M.; Profumo., A. Anal. Chem. 1993, 65, 2522-2527. (19) Szabadka, O.; Inczedy, J. J. Chromatogr. 1980, 201, 59-66. (20) Riley, J. P.; Taylor, D. Anal. Chim. Acta 1968, 40, 479-485. (21) Ceo, R. N.; Kazerouni, M. R.; Rengan, K. J. Radioanal. Nucl. Chem. 1993, 172 (1), 43-48. (22) Pasavento, M.; Biesuz, R. Anal. Chem. 1995, 67, 3558-3563. (23) Bruno, T. J.; Svoronos, P. D. N. CRC Handbook of Basic Tables for Chemical Analysis; CRC Press: Boca Raton, FL, 1989. (24) Shafer, M. M.; Overdier, J. T.; Hurley, J. P.; Armstrong, D. E.; Webb, D. Chem. Geol. 1997, 136, 71-97. (25) Lehto, J.; Paajanen, A.; Harjula, R.; Leinonen, H. React. Polym. 1994, 23, 135-140. (26) Martell, A. E.; Smith, R. M. Critically Selected Stability Constants of Metal Complexes Database. National Institute of Standards and Testing Standard Reference Database 46, Version 2.0; 1995. (27) Herrin, R. T. Ph.D. Dissertation, The University of Wisconsin, Madison, WI, 1999. (28) Bloom, N. S.; Crecelius, E. A. Anal. Chim. Acta 1984, 156, 139145. (29) Boyle, E. A.; Edmond, J. M. Anal. Chim. Acta 1977, 91, 189-197.
Received for review July 21, 2000. Revised manuscript received February 15, 2001. Accepted February 20, 2001. ES001509X
