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a Dep. of Plant and Soil Sciences, Univ. of Delaware, 152 Townsend Hall, Newark, DE 19717
b Dep. of Crop and Soil, Environmental Science, 236 Smyth Hall, Virginia Polytech. Inst. & State Univ., Blacksburg, VA 24061
c Dep. of Plants, Soils, and Biometeorology, College of Agriculture, Utah State Univ., Logan, UT 84322
* Corresponding author (eick{at}vt.edu)
Received for publication January 19, 2001.
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ABSTRACT |
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 TOP ABSTRACT INTRODUCTIONMATERIALS AND METHODSRESULTSDISCUSSIONCONCLUSIONREFERENCES |
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HAp, yet As(III) adsorption on ferrihydrite was greater than on goethite. The observed differences between this study and the earlier study on goethite are believed to be an intricate function of ferrihydrite's surface characteristics, which affect the mechanisms of adsorption and hence the affinity of organic acids such as HAp, FA, and CA for the ferrihydrite surface. As such, the adsorption of DOCs to ferrihydrite are assumed to be less favorable and to occur with a fewer number of ligands, resulting in lower surface coverage of weaker bond strength.
Abbreviations: AHS, acidified humic substance • CA, citric acid • DOC, dissolved organic carbon • DOM, dissolved organic matter • FA, Suwannee River fulvic acid • HAp, peat humic acid
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INTRODUCTION |
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TOPABSTRACTINTRODUCTIONMATERIALS AND METHODSRESULTSDISCUSSIONCONCLUSIONREFERENCES |
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Recently, we have demonstrated that As(V) and As(III) adsorption on goethite is decreased in the presence of humic acid peat (HAp), Suwannee River fulvic acid (FA), and citric acid (CA) (Grafe et al., 2001). The adsorption of As(V) on goethite was decreased in the order of HAp > FA > CA, and that of As(III) in the order of CA > FA > HAp. The decrease in As adsorption in the presence of DOC may be a function of the surface activity of the functional groups present on the DOC material. For example, humic acid peat, whose functional group content shows a higher phenol OH group content, reduced As(V) adsorption more than citric acid, whose functional group content is composed of three COOH groups as well as one OH group. In contrast, As(III) adsorption was greatly reduced by CA at lower pH values as the As(III) affinity for the surface decreased, and that of CA increased. Kaiser et al. (1997) demonstrated that the adsorption of DOC materials on ferrihydrite may be more physical in nature, while DOC adsorption on goethite was more chemical in nature. Hence, the overall weaker adsorption of DOC materials on ferrihydrite probably influences its competitiveness for surface sites with As(III) and As(V).
Unlike goethite, ferrihydrite is a poorly crystalline iron oxide whose random assembly of primarily dioctahedral Fe-octahedra results in an increased number of A- and C-type functional groups (Manceau, 1995; Waychunas et al., 1993). Several researchers have pointed out that As(V) and As(III) adsorption will take place preferentially on A-type hydroxyl groups of iron oxides (Sun and Doner, 1996; Manceau, 1995; Waychunas et al., 1993). Moreover, As(V) and As(III) differ in their use of adjacent B- and C-type hydroxyls. Arsenite binds preferentially with doubly coordinated C-type hydroxyls, while As(V) binds with triply coordinated B-type hydroxyls (Sun and Doner, 1996). The surface structural differences between goethite and ferrihydrite therefore warrant a closer examination of their effect on competitive adsorption reactions between As and other competing ligands such as organic acids.
The objectives of this research were to determine if dissolved organic acids decreased As(III) and As(V) adsorption on a poorly crystalline iron oxide (i.e., ferrihydrite) over a pH range of 3 to 11. Second, we examined the influence of DOC on As(V) transport with columns packed with ferrihydrite-coated sand previously aged with a leonardite humic acid.
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MATERIALS AND METHODS |
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TOPABSTRACTINTRODUCTIONMATERIALS AND METHODSRESULTSDISCUSSIONCONCLUSIONREFERENCES |
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The specific surface area was 253 m2 g-1, as determined by a five-point Brunauer–Emmett–Teller (BET) N2-gas adsorption isotherm method. X-ray diffraction (XRD) analysis and thermogravimetric analysis (TGA) were consistent with data presented in Schwertmann and Cornell (1991)(p. 61–80), verifying the identity of two-line ferrihydrite and showing no traces of either hematite or goethite.
Ferrihydrite coated sand was prepared for the column study with sand from T.J. Baker (Phillipsburg, NJ) (LOT H37720). The sand was washed with 0.1 M HCl, rinsed with double deionized water, and dried prior to coating. Four 100-g samples of previously cleaned and dried sand were measured into 250-mL ceramic evaporation dishes. Each 100-g sand sample received 8.08 g of Fe(NO3)3·9H2O and 50 mL of double deionized water. Following that, 12 mL of 5 M NaOH were added to each suspension and mixed well. Suspension pH was adjusted to 7.5 with appropriate amounts of either 5 M NaOH or 6 M HCl. The sand–ferrihydrite mixture was then allowed to dry in an oven at 313 K overnight. After drying, the sand was rewetted with double deionized water, and dried again in the oven at 313 K. After the sand was completely dry, the sand was transferred to a 53-µm stainless steel sieve and washed until all salts and indiscrete particles were removed. This point was determined when the effluent was a colorless, clear liquid. The coated sand was then dried again in the oven at 313 K (Amacher, 1999).
Adsorption Edges
Adsorption edges were measured to examine the function of pH on As and DOC sorption on two-line ferrihydrite. The adsorption of As(V) and As(III) and DOCs was examined as a function of pH (3–11) in a background electrolyte solution of 0.01 M NaNO3 at constant adsorptive (1.0 mM) and adsorbent concentrations (1.0 g ferrihydrite L-1). This adsorbent concentration was chosen in order to compare our results (on a surface area basis) with a previous study with goethite as the adsorbent (Grafe et al., 2001). An appropriate quantity of N2 gas–purged ferrihydrite suspended in 0.01 M NaNO3 was added to a 500-mL, Teflon-lined, flat-bottomed, water-jacketed reaction vessel and allowed to fully hydrate overnight. The reaction vessel was covered with a removable glass lid containing entry ports for a mechanical stirrer, pH electrode, N2 gas, burette tip, and a pipette. The pH was adjusted to 11.00 with a Brinkmann (Westbury, NY) Metrohm 718 stat titrino and the drop-wise addition of 0.10 M NaOH. Each ferrihydrite suspension adsorption edge was kept well stirred with the aid of a mechanical stirrer (Caframo [Ontario, Canada] RZR-2000) spinning at 300 rpm. All experiments were conducted at 298 ± 0.1 K and 0.101 MPa pressure under a N2 environment to eliminate CO2 influences.
After a minimum of 12 h, appropriate volumes of adsorptives were added to the suspension. Adsorptives for the adsorption edges were obtained from prepared stock solutions [0.10 M As(V) and 0.10 M As(III)] prepared from sodium salts. All DOC stock solutions were prepared with a background electrolyte of 0.01 M NaNO3. Suwannee River FA and HAp were obtained from the International Humic Substances Society (St. Paul, MN). An HAp solution of 0.01 M C was prepared by placing required amounts of dry humic acid in the background electrolyte at pH 8. The background electrolyte was previously purged with N2 gas for 20 min before the addition of the HAp. While dissolving, the pH was maintained at 8.00, and the system was continuously purged with N2 gas. The total acidity of the HAp is estimated according to reports from Stevenson (1994) and Navarro et al. (1993) to be approximately 7 mmolc g-1 of humic acid. Suwannee River FA stock solution was prepared (0.01 M C) in a similar manner at pH 7. The total acidity of Suwannee River FA has been reported to be 13.9 and 14.2 mmolc g-1 (Yates and von Wandruszka, 1999). A CA stock solution (0.01 M C) was prepared by dissolving an appropriate quantity of sodium citrate salt in 0.01 M NaNO3 at pH 7. The total acidity of CA can be calculated from its formula weight and the assumption that only COOH groups are active in the pH range of 3 to 11. The total acidity is then 15.62 mmolc g-1.
Adsorption edges were measured in duplicate and employed three consecutive addition scenarios for the adsorbates. The first scenario saw the addition of DOC species before the As species, the second scenario had the As species added before the DOC species, and the final scenario saw a simultaneous addition of the adsorptives. After the addition of each adsorptive, the pH was titrated back to 11.00 and allowed to equilibrate for a minimum of 2 h.
The suspension pH was lowered in half or full pH units from pH 11.00 to 3.00 with 0.10 M HNO3 with the aforementioned pH stat. After 2 h of minimum equilibration time, a sample was removed from the reaction vessel. For As(V), As(III), and DOC edges, 12.00-mL subsamples were removed with a Rainin (Emeryville, CA) automated digital pipette and filtered through a 0.10-µm Gelman (Ann Arbor, MI) metrical membrane into previously acid-washed polypropylene test tubes. The 12.00-mL subsample was divided into equal 6-mL samples for the analysis of As and total organic carbon (TOC). For the competitive adsorption edges, subsamples of 18.00 mL were taken for each adsorptive. The 18.00-mL subsample was divided again into equal volumes of 6 mL for As analysis, As(V) speciation, and TOC analysis. Arsenic was measured with a Spectro (Fitchburg, MA) inductively coupled plasma atomic emission spectrometer (ICP–AES). Samples containing any of the DOC materials were analyzed with a Tekmar–Dohrman (Mason, OH) Phoenix 8000 carbon analyzer.
Arsenic Speciation
Arsenate speciation was measured with a colorimetric assay (Cummings et al., 1999). Samples were mixed with appropriate amounts of 25 mM HCl and a reagent mix to give a final volume of 3 mL in a 4-mL plastic cuvette. The reagent mix consisted of four equal amounts of potassium antimony tartrate (0.544 g L-1), ammonium molybdate (24 g L-1), sulfuric acid (269.2 mL concentrated sulfuric acid L-1), and ascorbic acid (43.2 g L-1) solutions. The reagent mix was prepared fresh for every analysis within 2 h prior to the assay. Arsenate standards of 0, 250, 500, 750, and 1000 µM were prepared from stock solutions of known concentrations. Standards and samples were placed into a water bath of 351 ± 1 K for exactly 10 min, immediately removed, and placed in an ice bath for 5 min. Arsenate absorbance was measured with a Beckmann Coulter (Fullerton, CA) DU-640 spectrophotometer at a wavelength (
) of 640 nm. Any observed difference between As content established by inductively coupled plasma–atomic emission spectrometry and the colorimetric technique was attributed to As(III). No As(III) was detected. To determine potential interference from the organic acids, we prepared samples of known As(V) concentration in a background solution of the respective organic acid, and followed the procedure outlined above. No significant interference could be noted from the presence of HAp, FA, or CA.
Column Study
Two columns from Soil Measuring Systems (Tucson, AZ) were used. The inner diameter of the columns was 2.7 cm and the length of the columns was 7.6 cm. This gives each column a total volume of 43.55 cm3. Each column was uniformly packed with 67.2 g of ferrihydrite-coated sand, which resulted in a bulk density of the solid fraction of 1.54 g cm-3, and porosity equal to 0.42. One pore volume was thus equal to 18.29 cm3. Each column was attached to a fraction collector that was set up such that each test tube would collect 130 drops of solution, which corresponds to 8 mL. All solutions were passed through the columns at a flow rate of 0.4775 mL min-1. All solutions passing through the columns were at a pH of 6.5, and were prepared in a background solution of 0.01 M CaCl2. Columns were primed overnight by saturating the sand with 0.01 M CaCl2 prior to applying any of the adsorptives. The first column then received 18 pore volumes of 1.0 mM C leonardite humic acid solution (IHSS) followed up by enough CaCl2 (8 pore volumes until breakthrough of the DOC solution was achieved). Subsequently, we passed 1.0 mM As(V) solution through the column, followed by enough CaCl2 until breakthrough of the As(V) occurred. For the second column (after priming with 0.01 M CaCl2), we passed a 0.84 mM KBr solution through the column followed by enough CaCl2 solution until breakthrough of the KBr solution occurred. Potassium bromide serves as an inert background electrolyte or conservative tracer. Finally, four pore volumes of As(V) [1.0 mM As(V)] were passed through the column, followed up by enough CaCl2 to obtain breakthrough of As(V). We analyzed for Br- colorimetrically with a Lachat (Milwaukee, WI) QuickChem ion analyzer to develop a breakthrough curve (BTC) for Br-. Arsenic was measured by inductively coupled plasma–atomic emission spectrometry and leonardite humic acid was measured with a Tekmar–Dohrmann Phoenix 8000 carbon analyzer. The BTCs for As and leonardite humic acid were compared with that of Br- to determine the relative sorption of each species.
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RESULTS |
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TOPABSTRACTINTRODUCTIONMATERIALS AND METHODSRESULTSDISCUSSIONCONCLUSIONREFERENCES |
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Suwannee River fulvic acid adsorption on ferrihydrite is S-shaped from pH 3 to 11 with an adsorption maximum near pH 4. Similar adsorption behavior has been observed by Filius et al. (2000) for a soil fulvic acid on goethite at a surface loading of 300 mg L-1. Citric acid adsorption on ferrihydrite increases strongly from pH 11 to 9 and then remains nearly constant from pH 8 to 3. This adsorption behavior is different from the gradual adsorption increase on goethite as pH drops (Grafe et al., 2001; Geelhoed et al., 1998) and may be related to the increased density of reactive functional groups on ferrihydrite.
Kaiser et al. (1997) compared the adsorption of dissolved organic matter (DOM) and a fulvic and humic acid mixture (acidified humic substance [AHS]) on illite, gibbsite, goethite, and ferrihydrite, but saw no difference in the amount of adsorbed C kg-1 of goethite or ferrihydrite at equimolar loading rates. The overall sorption capacity of goethite was measured to be approximately 3 mol C kg-1 less than that of ferrihydrite, which indicated that the sorption capacity for DOM and AHS at the experimental loading rates for ferrihydrite had not been reached. Furthermore, this indicated that the surface density of DOM and AHS on goethite was therefore greater than on ferrihydrite.
Arsenate Adsorption in the Presence of Dissolved Organic Carbon
Only the addition scenario where DOC materials were added before As(V) is shown (Fig. 2)
, because there was no discernable difference between scenarios. Arsenate adsorption was inhibited by CA, while no effect was observed in the presence of HAp or FA. These observations are contrary to our parallel study with goethite, where HAp and FA inhibited As(V) adsorption on goethite, but where CA had no effect at the same adsorptive loading (1.00 mmol As or C 250 m-2; Grafe et al., 2001). Citric acid most effectively inhibited As(V) adsorption between pH 5 and 3 (17 to 20%, respectively). Geelhoed et al. (1998) observed a similar inhibition of phosphate adsorption on goethite in the presence of citrate.
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Dissolved Organic Carbon Adsorption in the Presence of Arsenite
Similar to the HAp–As(V) system (Fig. 3a), As(III) did not affect HAp adsorption (Fig. 5a)
. Increased HAp adsorption in the pH range of 3 to 5 should be regarded cautiously, because of coincidental precipitation reactions (Schulthess and Huang, 1991). It appears that the adsorption mechanisms of HAp and As(III) [and As(V)] are noninterfering.
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Similarly, As(III) inhibited CA adsorption on ferrihydrite by 17% from pH 10 to 6 (see Fig. 5c). From pH 10 to 8, CA adsorption increased and remained unaltered until pH 6 to 5, when CA adsorption increased again and reached an adsorption maximum near pH 4. From pH 7 to 3, the inhibition of As(III) adsorption on ferrihydrite by FA or CA occurred simultaneously with adsorption increases of the respective organic acid.
Column Study
The retention of the eluents on ferrihydrite-coated sand as a function of applied pore volumes followed the order of Br- < leonardite humic acid < As(V) As(V) after leonardite humic acid (Fig. 6)
. The graph clearly demonstrates that the greatest retention occurs for As(V) (approximately 12 pore volumes) regardless of previous treatment with leonardite (approximately 8 pore volumes). This is consistent with batch studies that demonstrated that HAp had no effect on As(V) adsorption on ferrihydrite.
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DISCUSSION |
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TOPABSTRACTINTRODUCTIONMATERIALS AND METHODSRESULTSDISCUSSIONCONCLUSIONREFERENCES |
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HAp, which is the reverse order of the results obtained with goethite. These results demonstrate the significance of the solid phase in adsorption and competitive adsorption processes between As and DOC species. The ability of CA to inhibit As(V) adsorption is generally in good agreement with the results of other researchers. Geelhoed et al. (1998) suggest that the adsorption of citrate to goethite is a combination of bidentate surface complexes of two COO- groups and hydrogen bonding by the remaining COO- group, which effectively competes with phosphate for surface sites at low pH. Similar results were obtained by other researchers who showed that relative functional group positioning on the organic acid and overall (small) molecular size increase competitiveness (Evanko and Dzombak, 1999; Struthers and Sieling, 1950). However, it remains unclear why CA was effective in inhibiting As(V) adsorption on ferrihydrite but not on goethite.
In contrast to CA, the fulvic and humic acids used in this experiment were not able to inhibit As(V) adsorption. This may be related to the functional group composition and the affinity and adsorption mechanism on ferrihydrite. Ferrihydrite is a poorly crystalline iron oxide whose random assembly of primarily dioctahedral Fe-octahedra results in an increased number of A- and C-type hydroxyls in comparison with crystalline iron oxyhydroxides (e.g., goethite) (Manceau, 1995; Waychunas et al., 1993). Sun and Doner (1996) and Waychunas et al. (1993)(1995) used a combination of spectroscopic techniques to examine the local bonding structure of As(V) on goethite and ferrihydrite surfaces. All researchers observed the formation of different surface complexes, which were dependent on surface coverage. Arsenate tetrahedra preferentially undergo ligand exchange reactions with A-type hydroxyls of Fe-octahedra, forming bidentate binuclear, bidentate mononuclear, and/or (at low molar As to Fe ratios) mononuclear surface complexes (Fendorf et al., 1997; Sun and Doner, 1996; Waychunas et al., 1993). Hydrogen bonding of a third oxygen with an adjacent B-type OH is more favorable for As(V) tetrahedra in binuclear coordination, as less strain is imposed on the Fe crystal (Sun and Doner, 1996; Waychunas et al., 1993). Therefore, the change in adsorbent phase from goethite to ferrihydrite favors the adsorption of As(III) and As(V).
Ferrihydrite also has a greater surface site density compared with goethite (16.8 vs. 5.73 µmol sites m-2; Dzombak, 1990). The surface site increase is reflected in two- to threefold adsorption increases of all adsorbates in this study on ferrihydrite relative to their adsorption on goethite (Grafe et al., 2001). Shape changes of HAp and FA adsorption edges are observed on goethite and ferrihydrite, while those of As(V) and As(III) remain similar for both adsorbents (Grafe et al., 2001). We believe that such a change in the shape of adsorption edges probably signifies a change in adsorption mechanisms or an alteration of functional group involvement in the adsorption process(es) of DOC materials.
Kaiser et al. (1997) investigated the adsorption of dissolved organic matter (DOM) and acidified humic substances (AHS, containing humic and fulvic substances) on goethite and ferrihydrite with 13C-nuclear magnetic resonance (13C-NMR) and diffuse reflectance Fourier-transform infrared (DRIFT) spectroscopy. The researchers demonstrated that the adsorption capacity for DOM and AHS were higher on ferrihydrite (7.1 and 7.8 mol C kg-1, respectively) than on goethite (4.0 and 4.5 mol C kg-1, respectively). Dissolved organic matter or AHS adsorption on goethite and ferrihydrite applied at equal adsorptive concentrations were the same regardless of adsorbent, which signified that DOM and AHS surface coverage was higher on goethite than on ferrihydrite, because suspension concentrations of goethite and ferrihydrite were the same (1 g to 1000 mL), while the specific surface area was about 4:1 in favor of ferrihydrite. Moreover, the intensity of DRIFT bands of protonated carboxyl groups in the spectra of DOM and AHS sorbed onto goethite indicated that the number of ligands involved in the adsorption reaction was higher on goethite compared with ferrihydrite. This occurred despite a closer arrangement of DOM or AHS molecules on the goethite surface. The authors concluded that the formation of polydentate complexes was favored on goethite, but not on ferrihydrite (Kaiser et al., 1997). The researchers also demonstrated that the intensity increases of COOH bands indicated the importance of these functional groups in adsorption reactions on ferrihydrite as well as on goethite.
In contrast, an increase in the intensity of phenolic bands was observed on goethite but not ferrihydrite, indicating their importance in adsorption on goethite. Furthermore, a greater intensity of carbohydrate bands for reactions on ferrihydrite suggests that DOM and AHS adsorption processes on ferrihydrite preferentially involve these carbohydrate rather than phenolic groups (Kaiser et al., 1997). In a similar study employing surface complexation modeling (CD-MUSIC), Filius et al. (2000) proposed an adsorption mechanism for a soil fulvic acid on goethite. Their results suggest that fulvic acid adsorbs via a strong outer-sphere adsorption of phenolic OH groups via hydrogen bonding, where the hydrogen ion served as a proton bridge satisfying charge on goethite–O- and fulvate–O- functional groups. Due to this weak electrostatic interaction, one would expect that As(V) would outcompete FA for adsorption on iron-oxide surfaces. Moreover, in the presence of adsorbed As(V), which strongly reduces the point of zero charge of variably charged surfaces, ligand exchange reactions between COO- groups of FA and functional groups of the surface should therefore be even less favorable.
The results of Kaiser et al. (1997) and Filius et al. (2000) help to explain the observed differences of As(V) adsorption on goethite and ferrihydrite in the presence of all three DOC materials. First, DOC adsorption to ferrihydrite is less favored than to goethite, resulting in weaker surface associations. Second, of all organic functional groups, COO- groups are more involved in adsorption processes on goethite and ferrihydrite than phenolic groups (phenol-OH) and/or carbohydrate functional groups (carbohydrate-OH). Hence, As(V) adsorption on ferrihydrite is least favored in the presence of CA, whose functional group content is composed of three COO- groups. Third, the relative importance of phenol groups in the adsorption process to ferrihydrite and goethite explains why HAp, with a higher phenol and catechol functional group content, was effective in decreasing As(V) adsorption on goethite but not on ferrihydrite. Finally, the increased involvement of carbohydrate functional groups demonstrates increased physical interactions between organic acids and the ferrihydrite surface. Moreover, hydroxyl functional groups on carbohydrate structures occur in a saturated carbon setting, where the dissociation of the OH is thermodynamically unfavorable in the pH range studied (Stevenson, 1994). Hence, FA and HAp were ineffective in competing with As(V) for surface sites.
Our column study provides further evidence that humic acids weakly adsorb to ferrihydrite-coated surfaces, as prior aging with leonardite did not effect the number of pore volumes required to achieve As(V) breakthrough on a ferrihydrite-coated sand–packed column. A lack of or decreased involvement of phenol-OH in the adsorption process on ferrihydrite explains why HAp did not inhibit As(V) adsorption on ferrihydrite, but was very effective on goethite, where phenol–OH surface interactions are more prevalent (Grafe et al., 2001).
Arsenite
Arsenite adsorption on ferrihydrite is reduced in the pH range from 7 to 3 in the presence of FA and CA with a concurrent increase in FA and CA adsorption. Arsenite adsorption on goethite involves bidentate binuclear surface complexes preferentially with A-type OHs of the surface, and hydrogen bonding between the remaining functional group and a neighboring C-type OH on the adsorbent surface (Sun and Doner, 1996). Several researchers have pointed out that ligand exchange reactions between COO- groups of organic acids and surface OHs is one of the major adsorption mechanisms for organic acids such as fulvic and humic acids, as well as for citric acid (Filius et al., 2000; Evanko and Dzombak, 1999; Geelhoed et al., 1998; Kaiser et al., 1997; Varadachari et al., 1997; Gu et al., 1994; Fontes et al., 1992; Sibanda and Young, 1986; Parfitt et al., 1977). Exchange reactions of ligands will take place preferentially at their pKa values, because deprotonated functional groups are required for ligand exchange. Therefore, As(III) adsorption on variably charged surfaces will take place preferentially at pH 9.29 (pKa1 of H3AsO3; Rubinson and Rubinson, 1998) and COO- adsorption of organic acids will take place at lower pH values (pka1–3 of CA are 2.24, 4.49, and 6.93, respectively). Therefore, As(III) adsorption on ferrihydrite becomes less favorable as pH drops, while that of FA and CA increases. The above explains the observed decrease in As(III) adsorption on ferrihydrite in the presence of FA and CA. The decrease in As(III) adsorption in the presence of FA or CA, however, cannot be ascribed merely to the presence of COO- functional groups, because one should see a small effect with HAp. Abundance and the density of COO- functional groups as well as the relative position of functional groups to each other appear to play an equally important role in the formation of surface complexes and hence in competitive reactions for surface sites (Evanko and Dzombak, 1999; Gu et al., 1994). Evanko and Dzombak (1999) showed that in adsorption reactions of various low molecular weight organic acids on goethite, equilibrium constants were highest for compounds having adjacent carboxylic groups, lower for compounds with adjacent phenolic groups, and lowest for compounds with phenolic groups in the ortho position relative to a carboxylic group. This may hence explain the greater As(III) adsorption inhibition by CA compared with FA and a lack of effect by HAp.
We also observed a shift in the adsorption maximum of As(III) on ferrihydrite compared with goethite (pH 7 vs. 9, Fig. 1). Other researchers have observed similar results, which may be explained by the higher functional group density of ferrihydrite compared with goethite (Pierce and Moore, 1982). It may also explain why CA and FA were less effective in decreasing As(III) adsorption on ferrihydrite than on goethite. On goethite, As(III) adsorption was reduced by 10, 15, and 46% at pH 3 (for the presence of HAp, FA, and CA, respectively), while on ferrihydrite As(III) adsorption at pH 3 was inhibited only by 0, 9, and 13% (HAp, FA, and CA, respectively). Additional explanation for this decrease is given by Kaiser et al. (1997), who showed that DOM and AHS adsorption to ferrihydrite (rather than goethite) involves fewer ligand functional groups. Therefore, the adsorption strength of HAp, FA, and CA on ferrihydrite is expected to be lower compared with goethite and may therefore may not be able to compete as effectively with As(III).
Arsenite adsorption on ferrihydrite was also inhibited by FA and CA between pH 11 and 7. This is unusual, because in this pH range both the ferrihydrite surface and FA and CA should be mostly negatively charged and hence repulsion should prevail, while As(III) adsorption should be more favorable, as the pKa1 of As(III) is approximately 9.29. Considering the individual adsorption of CA and As(III) on ferrihydrite (Fig. 1), we observe approximately 1.2 to 1.7 times greater CA adsorption than As(III) adsorption. In the presence of CA, As(III) adsorption is approximately 1.2 times less than in the absence of CA (Fig. 4). Inhibition of As(III) adsorption between pH 11 and 8 by FA may be due to an increased affinity of FA for the surface by electrostatic attraction of protonated amine sites in the FA structure and the negatively charged ferrihydrite surface. Furthermore, carbohydrates partitioning to the surface may physically block adsorption sites on ferrihydrite (Kaiser et al., 1997).
Arsenite adsorption on ferrihydrite was not inhibited by HAp, and likewise HAp adsorption on ferrihydrite was not inhibited by As(III). This is similar to what we observed for the As(V)–HAp system on ferrihydrite. However, this is in contrast to what was observed for the As(III)–HAp system on goethite (Grafe et al., 2001). We believe that the same reasons provided in the As(V)–HAp system are valid to explain why As(III) adsorption was not altered in the presence of HAp and vice versa; for example, increased physical partitioning of HAp through carbohydrate–OHs did not inhibit As(III) adsorption on ferrihydrite.
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CONCLUSION |
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TOPABSTRACTINTRODUCTIONMATERIALS AND METHODSRESULTSDISCUSSIONCONCLUSIONREFERENCES |
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The differences observed between goethite and ferrihydrite may be explained as follows:
1. Dissolved organic carbon (DOC) adsorption on ferrihydrite is less favorable, and results in a lower surface coverage with fewer ligands associating with the surface, congruent with findings of Kaiser et al. (1997). The bond strength between DOCs (FA and HAp) and the ferrihydrite surface is hence probably weaker.
2. Of all DOC functional groups, COOH groups are likely to be the most involved in the adsorption process. Moreover, the relative density and positioning of these functional groups on the organic acid as well as their size are important factors in competitive adsorption processes, as has been observed earlier (Evanko and Dzombak, 1999; Kaiser et al., 1997; Struthers and Sieling, 1950). Hence we observe an effect on As(V) adsorption on ferrihydrite in the presence of CA but not with FA.
3. The affinity constant of ligand exchange reactions between the fulvate molecule and the goethite surface according to the CD-MUSIC model was shown to be low (Filius et al., 2000). Furthermore, ligand exchange reactions are caused by strong electrostatic attractions between the fulvate molecule and the positively charged surface OHs (Filius et al., 2000). Arsenate, though, is known to lower the point of zero charge of variably charged surfaces, hence making ligand exchange reactions between the fulvate molecule and the ferrihydrite surface less likely.
4. Peat humic acid had no effect on the adsorption of As(V), because phenol groups are not likely to interact with a ferrihydrite surface as much as they do with a goethite surface (Kaiser et al., 1997).
5. Associations between the ferrihydrite surface and DOC involve more carbohydrate OHs forming weaker hydrogen bonds and other physical interactions with the ferrihydrite surface, and are therefore inhibited by As(V) (Kaiser et al., 1997).
6. Citric acid and FA decreased As(III) adsorption to a smaller extent, due to a greater surface affinity of As(III) with decreasing pH, but also for reasons mentioned above under Points 2 and 4.
7. We observed no effect of HAp on As(V) or As(III) adsorption and vice versa. We believe that the adsorption processes of As and HAp on ferrihydrite are independent and may not interfere with each other.
It has been demonstrated that the ability of DOC (HAp, FA, or CA) to inhibit As adsorption is in part a function of the adsorbent phase. Very little is known about the strength of adsorption of soluble organic matter on various geologic materials; however, the adsorption strength and the underlying mechanisms of adsorption may control the potential bioavailability of As in environments of relatively high DOC concentrations. Moreover, DOC may influence the potential bioavailability of other oxyanions such as phosphate, molybdate, and sulfate. These results demonstrate the importance of naturally occurring ligands in influencing the potential bioavailability of oxyanions in natural systems. Ferrihydrite, representative of a poorly crystalline iron oxide, shows promising characteristics to be a remediation tool in water treatment applications for As and other oxyanions in the presence of DOCs. Its applicability, however, requires short- as well as longterm kinetic studies to determine its remediation potential.
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ACKNOWLEDGMENTS |
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  M. E. Essington and R. M. Anderson Competitive Adsorption of 2-Ketogluconate and Inorganic Ligands onto Gibbsite and Kaolinite Soil Sci. Soc. Am. J., May 1, 2008; 72(3): 595 - 604. [Abstract] [Full Text] [PDF]
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M. E. Essington, J. B. Nelson, and W. L. Holden Gibbsite and Goethite Solubility: The Influence of 2-Ketogluconate and Citrate Soil Sci. Soc. Am. J., June 2, 2005; 69(4): 996 - 1008. [Abstract] [Full Text] [PDF]
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 E. Lombi, R. E. Hamon, G. Wieshammer, M. J. McLaughlin, and S. P. McGrath Assessment of the Use of Industrial By-Products to Remediate a Copper- and Arsenic-Contaminated Soil J. Environ. Qual., May 1, 2004; 33(3): 902 - 910. [Abstract] [Full Text] [PDF]
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