Environ. Sci. Technol. 1999, 33, 2252-2256
Interaction of Metal Ions with Acid Sites of Biosorbents Peat Moss and Vaucheria and Model Substances Alginic and Humic Acids† RAY H. CRIST AND J. ROBERT MARTIN Messiah College, Grantham, Pennsylvania 17027 DELANSON R. CRIST* Department of Chemistry, Georgetown University, Washington, D.C. 20057
The interaction between added metal ions and acid sites of two biosorbents, peat moss and the alga Vaucheria, was studied. Results were interpreted in terms of two model substances, alginic acid, a copolymer of guluronic and mannuronic acids present in marine algae, and humic acid in peat moss. For peat moss and Vaucheria at pH 4-6, two protons were displaced per Cd sorbed, after correction for sorbed metals also displaced by the heavy metal. The frequent neglect of exchange of heavy metals for metals either sorbed on the native material or added for pH adjustment leads to erroneous conclusions about proton displacement stoichiometry. Proton displacement constants KexH decreased logarithmically with pH and had similar slopes for alginic acid and biosorbents. This pH effect was interpreted as an electrostatic effect of increasing anionic charge making proton removal less favorable. The maximum number of exchangeable acid sites (capacity CH) decreased with pH for alginic acid but increased with pH for biosorbents. Consistent with titration behavior, this difference was explained in terms of more weak acid sites in the biosorbents.
Introduction A continuing environmental problem is the presence of toxic metals in natural waters. Recent promising approaches have been reported (1) to decrease their concentration by sorption to various biosorbents such as algae (2-4), peat moss (5, 6), fungi (7), microbial cells (8), and chitosan (9). One of the main mechanisms of sorption is metal binding to a localized anion site by displacing either an existing metal (ion exchange) or a proton (proton displacement), depending on pH. In their acid forms, these sites include alginic acid (a copolymer of guluronic and mannuronic acids) and sulfate acid groups for marine algae (10-12), ammonium and phenolic groups for marine phytoplankton (13), carboxylic acid groups on polysaccharides in cell walls for freshwater algae (14), and carboxylic acid and phenolic groups in fulvic and humic acids present in peat substances (15). Metal-metal exchange has been quantified for humic materials (16), peat (17), peat moss (6), and algae (18) in an † Presented in part at the Engineering Foundation Conference, Big Sky, MT, August 1996. * Corresponding author phone: (202)687-1682; fax: (202)687-6209; e-mail:
[email protected].
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ion exchange relation, shown in eq 1 for exchange with Ca. An apparent equilibrium constant Kex was calculated from eq 2
M2+ + (CaX2) ) Ca2+ + (MX2) Kex )
[Ca2+](MX2) [M2+](CaX2)
(1) (2)
where ( ) indicates the amount of sorbed metal on the solid phase. While sorption is frequently treated (13, 19, 20) as adsorption in Langmuir, Freundlich, or Scatchard plots, it has been shown (18) that such plots of ion exchange systems have linear segments that can erroneously be interpreted as multiple binding sites. Metals also displace protons from humic substances (21), as concluded from a pH drop when divalent or trivalent metals are added to a suspension of the biosorbent (22). From this indirect evidence, a process analogous to ion exchange (eq 3) with corresponding equilibrium expression 4 was postulated (15, 16):
M2+ + 2(HX) ) (MX2) + 2H+ KexH )
[H+]2(MX2) [M2+](HX)2
(3) (4)
However, the number of protons per divalent metal has been measured as 2 (25). Values 2 have been ascribed to deprotonation of hydrated metal ions upon binding (15). For algae, the stoichiometry of eq 3 was found for Cu-displacing protons from Vaucheria (27) and for Cd and Pb from Saccharomyces fluitans (12). Of considerable importance for applications of toxic metal removal by biosorbents, it was found (15) that these proton displacement constants KexH decreased markedly with pH and in a logarithmic relationship for Cd and peat moss (6). We now report the results of experiments with other sorbents and metals to investigate the generality and nature of this pH effect. Alginic and humic acids are of particular interest since they are well recognized substances and the active components of marine algae and peat moss, respectively. Vaucheria and peat moss provided information on examples of biosorbents of current interest in detoxification. Since covalent binding of protons with anion sites is a major component of KexH, the pH titration behavior of sorbents was studied as well as the stoichiometry of the proton displacement process.
Experimental Section Alginic acid (Sigma), polygalacturonic acid (Sigma), and humic acid sodium salt (Aldrich) were used as obtained. Canadian sphagnum peat moss was obtained from Berger, Inc. (Quebec), who provided the following characterization: electrical conductivity, 0.05-0.10 mS; humic acid, 3.5-4.0%; ash content, 0.5-1.0%; organic matter, 97-99%. This material was processed with a Thomas tissue grinder to reduce large particles. A 100-mesh sample of known dry weight was added to water to give a fine particle suspension of 20 mg/10 mL that was transferred conveniently by pipet and had a pH of 10.1021/es981278i CCC: $18.00
1999 American Chemical Society Published on Web 05/18/1999
TABLE 1. Apparent and Actual Ratios of Displaced Protons to Cda As Calculated from Total Cd Sorbed (CdTOT) or Cd Sorbed in the Proton Process (CdH) offb
onc
ratio
sorbent
pH
H
Ca
CdTOT
CdHb
apparent, H/CdTOT
actual, H/CdH
peat moss
4.00 5.00 6.00
167 270 212
39 55 169
195 321 380
156 266 211
0.86 0.82 0.55
Vaucheria
4.00 5.00 6.00
50 135 145
61 72 172
112 220 283
51 148 111
0.45 0.61 0.55
1.07 1.01 1.01 av 1.03 ( 0.03 0.98 0.91 1.30 av 1.10 ( 0.20
a All amounts in µequiv g-1 sorbent. b H off is given by the amount of LiOH used to hold the pH constant when Cd is added to the suspension. Caoff is given by the difference between Ca in the solution initially and after equilibrium with Cd. c CdTOT is given by the difference between Cd added and that in the solution at equilibrium. It is also given by determination of the (CdX2) in the sample after its treatment with HNO3. Values for the two methods agreed to within 10%, and the average value is given here. CdH is given by CdTOT minus CaOFF. All analyses were by performed by AA.
4.5 ( 0.1. Vaucheria was harvested from local limestone spring waters and stored at 0 °C. The sample was clipped fine and then treated with a Thomas tissue grinder. The pH of a suspension of this native material was 8.0 ( 0.1. Metal ions were provided as metal nitrates, and distilled water used for solutions. Titration Procedure. For alginic acid, increasing amounts of 0.10 M LiOH were added to a suspension of 20 mg in 10 mL of water, with pH measured after each addition. For Vaucheria, the acid form was prepared by suspending a sample in a solution of nitric acid at pH 1.5 for 20 min and then washing with water. To each of seven vials was added 10 mL of a stock suspension of 20 mg/10 mL, followed by 0.1, 0.2, 0.4, 0.7, 1.0, 1.5, and 2.0 mL of 0.01 M Ca(OH)2, respectively. The vials were capped and stirred mechanically, and the pH was measured after 4 h. For peat moss, 0.2 g was suspended in 100 mL of water, and the titration procedure for Vaucheria was followed. For humic acid, 20 mg (sodium salt) in 10 mL was brought to pH 4.5 with nitric acid, and increasing amounts of 0.10 M LiOH were added. Stoichiometry. A 20-mg suspension of peat moss or the acid form of Vaucheria in 10 mL of water was brought to the desired pH by Ca(OH)2, and then 0.050 mL of 0.10 M Cd was added. The amount of 0.01 M LiOH needed to hold pH constant gave the protons displaced. The solution and the sample (after nitric acid treatment to desorb metals) were analyzed for Cd and Ca by AA. The amount of Cd sorbed was gotten as the difference between Cd added and Cd in the solution at equilibrium. The amount from a sample analysis agreed to within 10% (usually within 5%) and the average given in Table 1 for Cd sorbed. Proton Displacements: General Procedure. An experiment for Cd with peat moss at pH 6 illustrates the general procedure, with data given in Table 2. A 100-mg sample of peat moss was suspended in 50 mL of H2O and brought to pH 6.0 with 0.01 M Ca(OH)2, which necessarily increased the (CaX2) component of the sample and Ca in solution. One 10-mL aliquot was used as a blank to obtain Ca in the solution and in the sample. Cd was added to the other aliquots to give 0.5, 1.0, 2.0, and 4.0 mM initial aqueous concentrations. The amount of 0.01 M LiOH needed to hold pH constant gave the amount of proton released HOFF in each case. After equilibration, suspensions were vacuum filtered. The filtrate was analyzed for Ca and Cd, and the sample was analyzed for (CaX2) by treatment with HNO3. Ca released was the difference between its aqueous concentration after equilibration with Cd and its value from the blank. Total sorbed Cd, (CdX2) and referred to as CdTOT, was taken as the sum of Ca and protons released in equivalents. Unreacted (HX) needed for calculation of KexH by eq 4 is given by (HX) ) (HX)O - HOFF, where (HX)O is the initial
TABLE 2. Proton Displacement Experimenta experimentb 1 Cd, mM HOFF, µequiv g-1 CaOFF, µequiv g-1 (CdX2), µequivc g-1 (CaX2), µequiv g-1 [Cd], mequiv L-1 (HX), µequivd g-1 KexH × 106
2
3
0.5 1.0 2.0 150 260 325 180 270 400 330 530 725 325 240 150 0.32 0.88 2.35 350 242 175 8.4 10.4 10.0 av 9.3 ( 0.85 × 10-6
4 4.0 375 400 775 85 5.74 125 8.6
a Equilibrium amounts for various concentrations of Cd with peat moss at pH 6.00. b Each column is for 20 mg of peat moss in 10 mL of water, but data are expressed as g dry weight of peat moss/L. c (CdX2) ) HOFF + CaOFF. d (HX) ) (HX)O - HOFF where (HX)O ) 500 µequiv g-1.
amount of acid sites at pH 6. The value of (HX)O for a given pH can be estimated by extrapolation of a curve (not shown) of protons released vs Cd added. A better value of (HX)O was gotten by a minimization procedure in which the value of (HX)O used in the calculation was adjusted until KexH values calculated for a series with different initial concentrations of added metal have a minimum error. In the case of Li, the pH was adjusted with LiOH. However, since some (LiX) dispersed into the aqueous medium, the suspension was centrifuged, and Li in the solution was determined by AA. The pellet was resuspended in 10 mL of H2O. When LiCl was added, the amount of protons released was measured and KexH was calculated as described above. The experiment was repeated for two other concentrations of LiCl, and the constants were averaged. Other Sorbents. For alginic acid, a suspension of 20 mg of alginic acid in 10 mL of water was brought to the desired pH with LiOH. A solution of a given metal was then added in successive amounts to give concentrations from 1 to 6 mM. Proton release lowered the pH, which was then brought back with 0.1 M LiOH. The amount of LiOH needed gives protons released and also other quantities for calculation of KexH. The Li used for adjusting pH was not incorporated in the polymer as shown by a blank experiment in which a 20-mg suspension was brought to pH 5.0 by LiOH. On centrifugation, the polymer formed a gel at the bottom of the tube, and analysis of Li in the supernatant by AA showed essentially all of the Li in the aqueous phase. For Vaucheria, 20 mg of the acid form was suspended in 10 mL of H2O giving a pH of 4-4.5. The suspension was brought to the desired pH by addition of 0.010 M Ca(OH)2, and the general procedure was followed. VOL. 33, NO. 13, 1999 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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For humic acid, 20 mg (sodium salt) was added slowly with stirring to 10 mL of water followed by nitric acid to achieve the desired pH. When 0.05 mL of 0.10 M Ca was added, the pH was maintained by adding 0.01 M LiOH, giving the amount of protons released. After equilibration (ca. 10 min), successive amounts of Ca were added giving solutions of 0.5, 1.0, 2.0, and 4.0 mM in Ca. Since the sample could not be separated by filtration or centrifugation, (CaX2) was taken to be the amount of protons released. The value of (HX)O was determined as in the general procedure.
Results and Discussion Stoichiometry. The ratio of equivalents of protons displaced from peat moss and Vaucheria per equivalent of Cd sorbed has apparent values based on total sorbed Cd (CdTOT in Table 1). However, these ratios, which vary from 0.45 to 0.86, do not represent the stoichiometry of the proton displacement process itself since some of the sorbed Cd displaced Ca (or Ca and Mg for Vaucheria). To obtain the actual ratio for only proton displacement, one must compare the protons off with the quantity (CdON - CaOFF), since this is the amount of Cd sorbed in the proton displacement reaction (CdH). As shown by the last column in Table 1, this actual ratio is 1.0, which corresponds to 2 protons off per Cd sorbed in the proton reaction. Proton Displacement Experiment. Results for a typical set of experiments with peat moss are shown in Table 2 to illustrate the method used to calculate KexH. A sample of peat moss (20 mg) suspended in 10 mL of 0.05 mM Cd at pH 6 released 3.0 µequiv of protons (HOFF) based on the LiOH used to keep the pH at 6. Analysis of the solution gave [Ca] and of the sample gave (CaX2), while (CdX2) is the sum of H and Ca released. As described in the Experimental Section, the amount of acid sites remaining at equilibrium (HX) is the difference between (HX)O, the initial amount of acid sites (acid capacity CH at pH 6), and the amount of (HX) that reacted (HOFF). Values for KexH calculated from these data by eq 4 and similar data for experiments with increasing amounts of Cd were averaged (see Table 2). We treat the KexH as a constant over the given Cd concentration range and assume it varies less than 10% (the average deviation for the 8-fold concentration range of 0.025-0.200 M). The available pH range for study depends on the metal. To obtain a measurable displacement by the weaker binding Ca, a high pH is needed to make eq 3 more favorable by mass action. Capacity CH of Acidic Components of Biosorbents. There are several aspects relating to a metal proton reaction. One factor is the maximum amount of sites left at a given pH, as determined by proton removal by hydroxide. For alginic acid, a pH titration by LiOH gave a sigmoidal-type curve (Figure 1) characteristic of relatively strong acids of similar strength, in this case the carboxylic acid groups of guluronic and mannuronic acid components of the polymer. From the curve, the full capacity of a suspension of 20 mg/10 mL is about 1.0 mL of 0.10 M LiOH where the pH rises to ca. 11. At pH 4.0 where 0.4 mL of LiOH has been used, the amount left (or CH in terms of mL of LiOH) is 1.0-0.40 ) 0.60 mL or 60% of its total value, while at pH 5.00, it is 1.0-0.8 ) 0.2 mL or only 20% of the original total sites present. Similar considerations apply to titration curves for biosorbents in Figure 1, although they are drawn out due to the presence of groups with a range of acid strengths. For example, carboxylic acid and phenolic groups of fulvic and humic acids in peat have pKa values that vary from 2.5 to 11 (28). Another factor is that metals differ in ability to displace protons, as shown in Figure 2 for metals displacing protons from Vaucheria at pH 6. At any given metal concentration, the ability of proton displacement is Pb > Cd > Ca. At high 2254
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FIGURE 1. pH titration of alginic and humic acids with 0.10 M LiOH (A) and Vaucheria and peat moss with 0.01 M Ca(OH)2 (B) for 20 mg of sorbent in 10 mL of water.
FIGURE 2. Displacement of protons from Vaucheria by various metals at pH 6.
concentrations, these curves reach a saturation level which is the capacity for that metal at that pH. For Vaucheria, the capacity for both Pb and Cd increase with pH, opposite to the trend for alginic acid as shown in Figure 3. An explanation for this effect is that the nature of the acid group, or more specifically, its acid strength, is also a factor for capacity. At the higher pH, stronger acids have been ionized, and only weaker acids remain in their acid form suitable for the proton displacement reaction. It appears from Figure 3 that for Vaucheria there are more of the weaker
FIGURE 3. Dependence of capacity CH on pH for alginic acid (A) and Vaucheria (B) with various metals.
FIGURE 4. Displacement of protons from humic acid by Ca at various pH.
acid groups that become available only at higher pH. A similar interpretation can be used to explain the increased protons displaced from humic acid by Ca at higher pH (Figure 4). Variations in KexH with Metal. In some cases, values of KexH calculated by eq 4 were found to depend on the metal at a given pH (see Figure 5). For Vaucheria at pH 6.0, KexH was 3.3 × 10-3 for Cd and 3.2 × 10-5 for Pb or a factor of 100 lower. This is at first surprising since Pb seems to be more active on the basis of its CH by a factor of about 5 (see Figure 3). However, CH enters the expression for KexH as a squared term in the denominator, since (HX) ) (HX)O - HOFF, and (HX)O is approximated by CH. Squaring the larger capacity for Pb thereby reduces its KexH. One must therefore conclude that there is no apparent correlation between KexH and the binding strength of a metal. pH Dependence of KexH. All sorbents showed a similar logarithmic decrease of KexH with pH, as shown in Figure 5. One explanation for such a decrease is that more than two protons are displaced, as reported for Cu with acid peat (15) and Cu and Pb on humic acid (25). However, in the present case the stoichiometry was shown to involve only two protons. An explanation for the KexH decrease in the present study is that Ka for acid groups decrease with pH. Following Kadlec and Keolian (15), one can consider KexH as a combination of eqs 5 and 6, where the demonstrated stoichiometry requires
FIGURE 5. Dependence of log KexH on pH for displacement of protons from Vaucheria and alginic acid (A) and peat moss and humic acid (B). squaring Ka in eq 7:
(HX) ) H+ + XM2+ + 2X- ) (MX2)
Ka
(5)
Kf
(6)
KexH ) Ka2Kf
(7)
log KexH ) 2 log Ka + log Kf
(8)
For alginic acid, with only one type of acidic group, the decrease in KexH with pH can be interpreted as lower values of Ka for remaining acid groups due to anion charge buildup at high pH (29). In support of this interpretation, it was reported that log Ka for apparent acid constants vs pH had a slope ca. -0.7 for a Suwannee River fulvic acid (30), which may however have other factors involved (31). Substituting this relationship into eq 8 gives a slope for log KexH vs pH of VOL. 33, NO. 13, 1999 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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2 (-0.7) or -1.4, comparable to those in the present study. The slope for Li (Figure 5B) is less than those for divalent metals in accord with only one proton displaced and hence a coefficient of 1 for log Ka in eq 8. An additional factor is that at higher pH, weaker acids are the ones in acid form, and their lower Ka values would also lead to lower observed KexH.
Literature Cited (1) Volesky, B., Ed. Biosorption of Heavy Metals; CRC Press: Boca Raton, FL, 1990. (2) Crist, R. H.; Oberholser, K.; Shank, N.; Nguyen, M. Environ. Sci. Technol. 1981, 15, 1212-1217. (3) Darnell, D. W.; Greene, B.; Henzel, M. T.; Hosea, J. M.; McPherson, R. A.; Sneddon, J.; Alexander, M. D. Environ. Sci. Technol. 1986, 20, 206-208. (4) Leusch, A.; Holan, Z. R.; Volesky, B. J. Chem. Technol. Biotechnol. 1995, 62, 279-288 and references therein. (5) Jeffers, T. H.; Ferguson, C. R.; Bennett, P. G. In Mineral Bioprocessing; Smith, R. D., Misra, M., Eds.; The Minerals, Metals, and Materials Society: Warrendale, PA, 1991; pp 289-298. (6) Crist, R. H.; Martin, J. R.; Chonko, J.; Crist, D. R. Environ. Sci. Technol. 1996, 30, 2456-2461. (7) Guibal, E.; Roulph, C.; Le Cloirec, P. Environ. Sci. Technol. 1995, 29, 2496-2503. (8) Golab, Z.; Orlowska, B.; Smith, R. W. Water Air Soil Pollut. 1991, 60, 99-106. (9) Guibal, E.; Jansson-Charrier, M.; Saucedo, I.; Le Cloirec. Langmuir 1995, 11, 591-598. (10) Kloareg, B.; Quatrano, R. S. Oceanogr. Mar. Biol. Annu. Rev. 1988, 26, 259-315. (11) Crist, R. H.; Oberholser, K.; McGarrity, J.; Crist, D. R.; Johnson, J. K.; Brittsan, J. M. Environ. Sci. Technol. 1992, 26, 496-502. (12) Fourest, E.; Volesky, B. Environ. Sci. Technol. 1996, 30, 277282. (13) Gonzalez-Davila, M.; Santana-Casiano, J. M.; Perez-Pena, J. Environ. Sci. Technol. 1995, 29, 289-301.
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(14) Percival, E. Br. Phycol. J. 1979, 14, 103-117. (15) Kadlec, R. H.; Keolian, G. A. In Peat and Water; Fuchsman, C. H., Ed.; Elsevier: London, 1986; pp 61-93. (16) Sposito, G. In CRC Critical Reviews in Environmental Control; CRC Press: Boca Raton, FL, 1986; pp 193-229. (17) Baes, A. U.; Bloom, P. R. Soil Sci. 1988, 146, 67-72. (18) Crist, R. H.; Martin, J. R.; Carr, D.; Watson, J. R.; Clarke, H. J.; Crist, D. R. Environ. Sci. Technol. 1994, 28, 1859-1866. (19) Benedetti, M. F.; Milne, C. J.; Kinniburgh, D. G.; Van Riemsdijk, W. H.; Koopal, L. K. Environ. Sci. Technol. 1995, 29, 446-457. (20) Brownawell, B. J.; Chen, H.; Collier, J. M.; Westall, J. C. Environ. Sci. Technol. 1990, 24, 1234-1241. (21) Stevenson, F. J. In Micronutrients in Agriculture, 2nd ed.; Luxmoore, R. J., Ed.; Soil Science Society of America: Madison, WI, 1991; Chapter 6 and references therein. (22) Van Dijk, H. Geoderma 1971, 5, 53-67. (23) Gosset, T.; Trancart, J.-L.; Thevenot, D. R. Water Res. 1986, 20, 21-26. (24) Bunzl, K.; Schmidt, W.; Sansoni, B. J. Soil Sci. 1976, 27, 32-41. (25) Stevenson, F. J. Soil Sci. 1977, 123, 10-17. (26) Wolf, A.; Bunzl, Dietl, F.; Schmidt, W. F. Chemosphere 1977, 207-213. (27) Crist, R. H.; Martin, J. R.; Guptill, P. W.; Eslinger, J. M.; Crist, D. R. Environ. Sci. Technol. 1990, 24, 337-342. (28) Morel, F. M. M. Principles of Aquatic Chemistry; WileyInterscience: New York, 1983; pp 237-311. (29) McBride, M. B. Environmental Chemistry of Soils; Oxford University Press: New York, 1994; pp 107-109. (30) Ephraim, J.; Alegret, S.; Mathuthu, A.; Bicking, M.; Malcolm, R. L.; Marinsky, J. A. Environ. Sci. Technol. 1986, 20, 354-366. (31) Marinsky, J. A.; Ephraim, J. Environ. Sci. Technol. 1986, 20, 349354.
Received for review December 9, 1998. Revised manuscript received March 29, 1999. Accepted April 16, 1999. ES981278I
