Uptake of Metals on Peat Moss: An Ion-Exchange Process

Although badly scattered, their data suggest a linear relation between log and pH. ..... A plot of log Kex for PM vs log KMAc was linear as shown in F...
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Environ. Sci. Technol. 1996, 30, 2456-2461

Uptake of Metals on Peat Moss: An Ion-Exchange Process† RAY H. CRIST, J. ROBERT MARTIN, AND JOSEPH CHONKO Messiah College, Grantham, Pennsylvania 17027

DELANSON R. CRIST* Department of Chemistry, Georgetown University, Washington, D.C. 20057

Various biomaterials have shown promise as sorbants to remove heavy metals from water. Several advantages of peat moss for such applications include its abundance, low cost, and high metal capacity. A Ca-loaded column of peat moss was therefore studied with mixtures of metals. Metals bind to anionic sites H ) or by displacing protons from acidic groups (Kex existing metals from anionic sites at high pH (Kex). These ion-exchange equilibrium constants were determined in batch experiments by direct measurement of species in solution and sorbed on the solid phase. The same Kex values of Mg 0.342, Mn 0.862, Ca 1.00, Ni 1.42, Zn 1.88, Cd 2.82, Cu 9.97, and Pb 26.7 relative to Ca were found for a given metal alone or in the presence of a mixture, thus showing that the metals function independently. Under conditions employed for a Mg/Mn mixture, it was found that ion-exchange equilibria were maintained along the column due to very fast rates for metal-metal exchange as measured in a separate kinetic study. A linear relationH vs pH over 106 was interpreted as ship for log Kex due to metal binding to sites of different acid strengths.

Introduction Despite considerable work in recent years, the presence of toxic metals in groundwater remains a serious environmental problem. Innovative approaches using the sorption of metals on various types of biomass appear promising (1-3). Peat moss, whose high metal capacity also enhances its use as a growing medium, is abundant and inexpensive and has been used for this purpose (4-8). Metal uptake on peat is actually a process of ionexchange at acidic sites that have resulted from the humification process. Extensive work over the years with soil has greatly increased our knowledge of such metal reactions (9-11). For peat, metals can react with carboxylic and phenolic acid groups of fulvic and humic acids (12) to † Presented in part at the Engineering Foundation Pollution Prevention Conference, Palm Coast, FL, January 1995. * Corresponding author telephone: (202) 687-1682; fax: (202) 6876209; e-mail address: [email protected].

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release protons (proton displacement, with equilibrium H constant Kex ) or, at sufficiently high pH, with their anion sites to displace an existing metal (with ion-exchange constant Kex). By titrating acid peat in the presence of Cu metals, Kadlec H and Keoleian (13) determined values for Kex for Cu as defined in our work with algae (14, 15). Although badly scattered, their data suggest a linear relation between log H H Kex and pH. In an empirical approach, Kex was estimated for various transition metals by the titration of acid peat in the presence of the metal (16). Ion-exchange constants were determined for mixtures of the agriculturally important nutrients K/NH4, Ca/Mg, and Ca/K with sphagnum peat (17). Although several other ion-exchange studies have been done on components of peat, namely, humic and fulvic acids, these also rely on measurements of aqueous species rather than determining sorbed amounts directly (18). We now describe the results that show that peat moss (PM) is an efficient sorbant to separate a mixture of heavy metals in a PM column pretreated with Ca. To understand this column operation, experiments were done to find out if each metal interacted independently with the solid phase. For this purpose, ion-exchange constants for a given metal were determined with that metal alone and in mixtures of metals. Analyses were done by direct measurement on all species in solution or sorbed on the solid phase, and thus the constants obtained are more reliable than those previously reported in which the metal content on the solid is uncertain. Rates of metal displacement of Ca and H from PM were also measured, since such factors would be important in either column or batch operations of PM as a metal scavenger.

Experimental Section Canadian sphagnum PM was obtained from Berger, Inc (Quebec), who provided the following characterization: pH 3.3-4.4; 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 sample of known dry weight was added to aqueous solutions to give a fine particle suspension that was transferred conveniently by pipet. Metal ions were provided as metal nitrates, and distilled water was used for solutions. To determine sorbed metal, the sample was treated with 3 M nitric acid for 10 min and filtered. The filtrate was analyzed for metal concentration by atomic absorption (AA) with a Perkin-Elmer Model 2380 A instrument. Data were obtained in triplicate and given in micromoles per gram dry weight (µmol g-1). Various forms of PM were obtained by treating 35-mesh native peat moss (PM-N) with different solutions. The acid form (PM-H) was prepared by taking a PM suspension to pH 1.0 with nitric acid for 1 h followed by thorough decantation washing. A suspension of the product has a pH of about 4.5. The Ca form (PM-Ca) was made by adding PM-N to a solution 0.1 M in Ca. After 30 min with the pH held constant at 6.5 by LiOH additions, the suspension, 0.1 M in Ca, was adjusted to pH 7.0 by 0.01 M LiOH, filtered, thoroughly washed, paper-dried, air-equilibrated, and sieved (35 mesh). When this material was treated with other metals, including the most strongly binding Pb and Cu,

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TABLE 1

Aqueous Concentrations, Sorbed Amounts, and Calculated Kex for Ca-PM in the Presence of a Mixture of Cd and Zn experiment

[Ca], mM [Cd], mM [Zn], mM (CaX2), µmol g-1 (CdX2), µmol g-1 (ZnX2), µmol g-1 ∑(MX2)a Kex for Cd

1

2

3

4

5

0.085 0.0041 0.0067 375 43 43 461 2.38

0.1000 0.0088 0.0117 308 75 78 461 2.76

0.1250 0.0235 0.0340 220 130 111 461 3.14

0.1750 0.0940 0.1080 131 178 157 466 2.53

1.45

2.16

1.86

0.1656 0.0620 0.0760 164 156 141 461 2.53 av 2.66 ( 0.22 1.86 av 1.85 ( 0.16

Kex for Zn a

1.94

Ion-exchange capacity.

TABLE 2

Kex for Reaction of Metals with PM-Ca, Vaucheria, and Calcium Alginate PM-Ca Li K Mg Mn Ca Ni Zn Cd Cu Pb

Vaucheria

calcium alginate 0.007

0.342 0.840 1.00 1.42 1.88 2.82 9.97 26.7

0.003 0.47

0.098

1.00

1.00

0.92 1.04 2.62 2.04

0.54 1.97 16.2

there was negligible release of protons and other metals, indicating that essentially all of the exchangeable sites in PM-Ca contain Ca. The Zn form (PM-Zn) was made similarly. Metal and Acid Content of Native Stock. The metal content of PM-N was determined by treating a 50-mg sample suspended in 100 mL of H2O with HNO3 to give successively pH values of 5, 4, 3, 2, and 1. After coming to equilibrium at each pH, 15-20 min, a 5-mL sample of the whole suspension was removed for metal analysis of the solution by AA. The acidity of PM-N was investigated by a technique reported for weak acid ion exchangers (19). Various amounts of 0.1 M NaOH and 0.1 M NaCl solutions were added to nine screw-cap (Teflon-lined) vials each containing 0.10 g of PM-N in 10 mL of H2O. The vials were placed on a shaker for various lengths of time, and the pH was measured after the settling of PM. Solutions at pH 6.5 and above were brown, indicating some solubilization of PM as reported by others (20). Ion-Exchange Equilibrium Constants for Metals. Four suspensions were made, each containing 20 mg of PM-Ca, 20 mL of H2O at pH 6.0, and a given metal typically at 0.1, 0.20, 0.35, and 0.50 mM, respectively. Any small change in pH was adjusted with 0.01 M LiOH or HNO3. After 20-30 min with stirring, the suspension was filtered and the PM sample treated with 3 M HNO3 for 10 min to release sorbed metals and filtered. Both filtrates were analyzed for metals, giving metal sorbed and Ca displaced. For the reaction

M2+ + PM-Ca f Ca2+ + PM-M the ion-exchange equilibrium constant Kex or selectivity

coefficient for non-ideal cases (21, 22) was calculated from (15) (22)

Kex )

[Ca2+](MX2)

(1)

[M2+](CaX2)

where ( ) indicates the amount of sorbed metal on the solid phase. This same procedure was used for a mixture of metals. For example, for PM-Ca in the presence of both Cd and Zn, the metal concentrations for 20 mg of PM-Ca in 20 mL of water were each 0.1, 0.20, 0.35, and 0.50 mM, respectively. After equilibration, samples and solutions were analyzed for Ca, Cd, and Zn as shown for a typical set of experiments in Table 1. For more weakly bound metals (Mg), concentrations were larger to give sorptions sufficiently above experimental error to be reliable, while for strongly bound metals (Pb) metal concentrations were lower. The values given in Tables 2 and 3 are the averages of 12-15 experiments, since each set of four or five concentrations for a given suspension was done in triplicate. H Proton Displacement Constants Kex . A suspension of 20 mg of PM-H in 20 mL of H2O was adjusted to pH 7, for example, and centrifuged. The pellet was taken up in 10 mL of H2O, and after adjustment to pH 7, Cd was added to give a concentration of 2 mM. The LiOH then added to hold the pH constant was a measure of the proton release that occurred. After 3 min for equilibration, the suspension was vacuum filtered quickly to keep the Li in the solid phase. Filtrate and sample were analyzed for Cd (as well as Li). The value of (HX) was determined as described below. For the reaction

Cd2+ + 2(HX) f (CdX2) + 2H+ H Kex was calculated from H Kex )

[H+]2(CdX2)

(2)

[Cd2+](HX)2

Three additional runs were made to give a Cd range of 0.2-2.0 mM. The procedure was repeated for pH 4, 5, and 6. Reaction Rates for Cation Exchange. Rates of metals displacing protons, which required about 30 min for completion, were followed by measuring pH. About 20 mg of PM-N was suspended in 20 mL of water, and 20 mL of metal solution was added to give the desired final con-

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TABLE 3

Ion-Exchange Constants Kex of Metals with PM-Ca one metal two metals Cd/Zn Cu/Mg Cu/Zn Mg/Zn Mn/Cd Ni/Cd Pb/Mg Pb/Zn three metals Cd/Mg/Zn mixture av mixture av/one metal

Cd

Cu

2.82 ( 0.21

9.97 ( 1.71

2.32 ( 0.26

11.2 ( 1.7 13.6 ( 1.8

3.26 ( 0.17 2.70 ( 0.36

3.31 ( 0.07 2.89 ( 0.38 1.02

Mg 0.34 ( 0.17

9

0.84 ( 0.04

1.42

Pb 26.7

0.260 ( 0.014

0.253 ( 0.013 0.28 ( 0.35 0.807

ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 30, NO. 8, 1996

Zn 1.88 ( 0.11 1.66 ( 0.19

0.864 ( 0.023

0.340 ( 0.007

12.4 ( 1.2 1.24

Ni

0.251 ( 0.017

centration. After the addition of the metal solution, the pH was maintained at 6.0 by addition of 0.01 M LiOH. A plot of milliliters of LiOH vs time was made with (mL)∞ representing the total amount of (HX) available for reaction at the start. The amount of (HX) at any time is given by (HX)t ) (HX)o - (HX)used ) (HX)o - LiOHt. Data were plotted according to a first-order rate equation ln Ct vs t, where Ct ) (HX)t, with slope of -k, the first-order rate constant. Rates of protons displacing metals from PM-M were studied by quickly adding enough HNO3 to the PM suspension at pH 6 to bring the pH to 4.0 (taken as t ) 0). As protons were taken up, additional HNO3 was added to maintain the pH at 4.0, and first-order plots were made for milliliters of H+ added vs time as described above. For rates of metal-metal exchange, 50 mg of PM-Ca was suspended in 50 mL of H2O and with rapid stirring mixed with 50 mL of 0.01 M solution of another metal. A 5-mL sample of the whole suspension was removed with a small dipper and vacuum filtered as rapidly as possible. Filtrates were analyzed for Ca displaced and rate constants obtained from plots of ln(CaX2) vs t. Metal Sorption within Column A. For qualitative experiments to show metal holdup along a column during elution, several 20-25-mesh samples of PM were converted to PM-Ca and placed in 1-in. sections of 4-mm glass tubing connected with vinyl tubing. The sections were separated at the joints with glass wool into which syringe needles were inserted for liquid sampling. A solution 1.0 mM in various metals at pH 6.0 was put through under slight positive air pressure to give a flow rate of 1 drop/s to ensure equilibrium. After about 2 h, samples were removed at the same rate from a given port (other ports stoppered) and analyzed. Equilibrium within Sorption Column B. To determine if ion-exchange equilibria are maintained within a column, a six-unit column was made consisting of alternate sections of 1.8 cm long × 2 mm i.d. glass tubing for the PM sample and 2.5 cm long × 7 mm i.d. plastic tube for the solution joined with vinyl tubing and the pieces separated by glass wool (see Figure 1). A 10-mg sample of PM-Ca (35-100 mesh) was put into each glass tube section. Breathing on the dry powder helps relieve static. A solution 1.0 mM in Mn and 0.30 mM in Ca at pH 6.0 was put through at 0.05 mL/30 s, and the effluent was analyzed for Mn breakthrough. When Mn was ca. 0.03 mM, the solutions and samples were obtained successively from the sections by

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Mn

0.864 1.02

1.40 ( 0.32 1.35 ( 0.07 1.44 ( 0.19

1.44 1.01

25.6 ( 1.2 24.5 ( 0.33

1.50 + 0.22

25.0 0.94

2.36 ( 0.07 1.65 0.88

FIGURE 1. PM-Ca Column for determining Kex at various points for Mn displacing Ca.

cutting the plastic connectors with a knife. The solution was taken out with a syringe while the sample was pushed out with a wooden stem. The solutions and the PM samples after a 10-min treatment with 3 M HNO3, were analyzed for metals.

Results and Discussion Metal Content and Acidity of Native PM (PM-N). Acidification of PM-N released metals as shown in Figure 2. The most abundant metal was Ca at 140 µequiv g-1 while Mg, K, Na, and Zn total 103 µequiv g-1. A titration curve of PM-N is shown in Figure 3, which also includes reported values (19) for Amberlite IRC-50, a

FIGURE 2. Protons displacing metals from native peat moss (PM-N) on acidification.

FIGURE 4. Relationship between ion-exchange constants Kex and formation constants for the corresponding metal acetates KMAc.

FIGURE 3. Titration curve for PM-N for samples after shaking with various amounts of aqueous NaOH after 16 h and after 8 days. For comparison, data from ref 10 on Amberlite IRC-50, a polymer of methacrylic acid and divinylbenzene, and a phenolic resin are also included.

polymer of methacrylic acid and divinylbenzene, and a phenolic resin. The shape of the curve with no evidence of an inflection point suggests a mixture of acid groups with a wide range of pKa values that may depend on pH (23). In contrast, isolated humic acids do show the usual sigmoidal curve (10) (12), although drawn out compared to salicylic acid (12), a model compound. The observation that protons are relatively easily displaced from PM, e.g., by Ca and Mg, may be due to the presence of relatively strong acids (12). It may also result from the chelation effect (24, 25), which would release protons from a carboxylic acid group when a divalent metal ion also binds to an adjacent phenolic OH or an adjacent carboxylic acid. Bidentate complex formation of Ca with carbohydrates, some with carboxylate groups, is a wellknown phenomenon responsible for important physiological activity of Ca in biological systems (26). The facile proton loss by metals has two important aspects: (1) for metal accumulation applications of PM, pretreatment with Ca solution would be very beneficial to prevent protons displaced from an early stage of the column from adding to anion sites thereby reducing ion-exchange capacity in the later stages and (2) for all quantitative work with PM-N or PM-H in this work, the pH was maintained

by adding LiOH and Kex values determined on PM-Ca samples, which do not have exchangeable (HX) sites. Ion-Exchange Constants. Values for Kex for various metals displacing Ca are given in Table 2. Also shown are reported values (15) for the alga Vaucheria and calcium alginate for reference, since alginic acid, a copolymer of manuronic and gluconic acids, is a component of marine algae (27). Because Kex is dimensionless, it is possible to compare the values for such different solids directly. The constants for PM are larger, especially for Pb, and in good agreement with values for peat-Ca with Cu, Ni, or Mn determined by an indirect titration method (16). It was previously found for algae (15) that Kex could not be explained entirely by ionic bonding but was also related to the covalent bond character of metal carboxylate as represented by the formation constant of metal acetates KMAc. A plot of log Kex for PM vs log KMAc was linear as shown in Figure 4, indicating that the free energies of the exchange process on PM correspond to metal carboxylate formation (28). Values of Kex for metals displacing Ca, alone and in the presence of other metals, were calculated from eq 1 using aqueous molar concentrations of metal ions and metal content on samples, assuming each metal system independent of the other. Typical data are shown in Table 1 for a Cd/Zn mixture. Aqueous concentrations varied in this case by a factor of 23 for Cd (with loadings from 9% to 38%) and of 16 for Zn (with loadings from 9% to 38%). For this wide range of concentrations, calculated Kex varied by only 8.2% for Cd and 8.6% for Zn. In general, initial concentrations were chosen so as to give reliable readings of aqueous and sorbed amounts, and loadings averaged between 10% (Mg) and 90% (Pb). Constants for each metal in several mixtures are averaged, and the ratio of the average in the various mixtures to the single metal is shown in the last row of Table 3. Ratios close to one, even for the three-component mixture, indicate that metal ions in solution act with the PM sites

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H FIGURE 5. pH dependence of Kex for Cd displacing protons from PM-H.

independently of other metals in solution or sorbed on the solid in the loading ranges studied. The effect of a pH change on Cd displacement of protons from PM-H was examined over a pH range of 4-7. As H shown in Figure 5, a plot of log Kex vs pH was linear over 6 10 . One interpretation of this result is that at lower pH stronger (HX) acid sites are present and Cd displaces protons H relatively easily leading to a large Kex . At higher pH, on the other hand, only weak acid sites are still in the (HX) form, and Cd displacing such protons is more difficult causing H a low Kex . A range of acid strengths of PM was in fact indicated by the non-sigmoidal titration curve of Figure 3 and the reported range of effective pKa of 2.5-11 for humic and fulvic acids (12). It should be noted that the present approach to determining ion-exchange constants is to measure the amounts of all species, sorbed as well as in solution. Since pH in Figure 5 was determined by adding LiOH, the measured amounts of [Li+], [H+], (LiX), and (HX) allowed H calculation of Kex for the reaction

Li+ + (HX) f (LiX) + H+ The value of this constant at pH 7, for example, was the same as that found for the direct reaction of Li+ with PM-H (29). Such a calculation is justified, since we have shown specifically (Table 3) that ions in mixtures function independently with the solid phase. Rates of Cation Displacement. A first-order rate plot for Ca displacing protons from PM-N at pH 6 is shown in Figure 6. After a very fast process, with k1 estimated as 260 × 10-4 s-1 from two points, there appears to be two distinct straight lines for slower processes with k2 ) 52.5 × 10-4 s-1 and k3 ) 17.3 × 10-4 s-1. Values of the slower rate constants were found to be independent of [Ca2+] over a range of 0.001-0.05 M. Averages of these constants are given in Table 4 where it can be noted that values of k2 are 2-15 times larger than k3. Similar plots were found for protons displacing Ca from PM-Ca and Zn from PM-Zn. As shown in Table 4, rate constants for displacing Ca are independent of [H+] over a concentration change by a factor of 2.5 (pH 4.1 vs pH 4.5). A detailed interpretation of these rates is difficult to give at the present time. The observation of k2 and k3 rates independent of aqueous concentrations implies slow ion transport within the solid phase. Factors influencing this process include ion-charge density, degree of hydration,

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FIGURE 6. First-order rate plot for Ca displacing protons from native peat moss (PM-N) by the reaction Ca2+ + 2(HX) f 2H+ + (CaX2) where Ct is the amount of (HX) sites remaining at time t (see text). TABLE 4

First-Order Rate Constants for Cation Displacements k1 × 104 (s-1)

k2 × 104(s-1)

k3 × 104 (s-1)

Ca, PM-N Ca, PM-H Cu, PM-N

Metal Displacing Protons 196 ( 13 43 ( 2 220 57 200 31 ( 1

15 ( 3 4 7(1

PM-Ca, pH 4.1 PM-Ca, pH 4.5 PM-Zn, pH 4.1

Protons Displacing Metals 60 ( 27 39 ( 1.3 60 ( 3 34 ( 2 107 22

17 ( 2.0 14 ( 4 11

PM-Ca, Zn

Metal-Metal Displacement 1866 1027

and microstructure of the sorbant. Of considerable importance for use of PM as a biosorbant, however, is the very fast metal-metal rate for Zn displacing Ca with a rate constant over eight times those of proton-metal reactions. Exchange Reactions in PM Columns. Preliminary experiments with native PM showed that metals easily displace protons from (HX) sites, causing the pH to decrease to about 2.5. This acid eluant protonates sites in later stages, thereby lowering the ion-exchange capacity to ca. 30% (see Figure 2). The present results were therefore obtained from columns of PM-Ca where (HX) sites had been exchanged for Ca. Concentrations of Mg/Zn along a PM-Ca column are shown in Figure 7. These results are for column A after several hours using a feed stock 1.0 mM in Mg and Zn. The stronger binding Zn (Kex ) 1.88 vs 0.342 for Mg) is preferentially held up in earlier sections leading to lower solution concentrations. A similar pattern of results was found for a mixture of five metals with those having higher Kex giving lower solution concentrations (see Figure 8). The actual system is complex, with advancing metals displacing sorbed metals according to relative Kex.

0.245 ( 0.56 in good agreement with that of 0.276 in Table 3. The present results clearly show that ion-exchange equilibria are maintained through the column under the conditions studied. Suitable data for the other more strongly bound metals could not be obtained, probably due to an initial, much more rapid decrease in concentration (see Figure 8). Metal concentrations for this column had to be much higher than found in practice in order to obtain suitable data for verifying Kex.

Acknowledgments We would like to acknowledge Berger, Inc. for providing the peat moss sample used and to thank Daniel Dupuis for his generous assistance. FIGURE 7. Concentrations of metals along a PM-Ca column in a flowing system 1.0 mM in Mg and Zn (column A).

FIGURE 8. Concentrations of metals along a PM-Ca column in a flowing system 1.0 mM in various metals (column A). TABLE 5

Ion-Exchange Equilibration along a PM-Ca Column with a Mn Feed Solutiona column section

concnb

1

2

3

4

5

6

[Ca2+]

0.34 0.37 0.42 0.54 0.72 0.80 0.94 0.88 0.82 0.60 0.31 0.12 [Mn2+] 350 333 228 134 47 (MnX2) 373 (CaX2) 162 179 190 278 308 382 Kex 0.832 0.819 0.896 0.738 1.01 0.819 a Feed solution 1.0 mM in Mn and 0.30 mM in Ca. b Aqueous concentrations [ ] in mM and sorbed amounts ( ) in µmol g-1.

To determine whether exchange equilibria are maintained along a column, Mn displacing Ca from PM-Ca was studied in detail using a feed stock of 1.0 mM Mn and 0.3 mM Ca with column B. The concentration of Mn before and after the first section of PM was 1.0 and 0.89 mM, respectively. The average, 0.94 mM shown in Table 5, is taken to represent the concentration of Mn in the first section. Sorbed Mn for this section, 373 µmol g-1, was determined by analysis of the entire amount of solid in that section. Other concentrations and amounts were determined similarly and are given in Table 5. Values of Kex were reasonably constant over all sections with an average of 0.852 ( 0.046, which compares very well with the value of 0.864 in Table 3. A similar experiment with Mg gave Kex

Literature Cited (1) Volesky, B., Ed. Biosorption of Heavy Metals; CRC Press, Inc.: Boca Raton, FL, 1990. (2) Darnall, D. W.; Greene, B.; Hosea, M.; McPherson, R.; Henzl, M. Alexander, M. D. In Trace Metal Recovery from Aqueous Solutions; Thompson, R. T.; Ed.; Burlington House: London, U.K., 1986; pp 1-24. (3) Lembi, C. A.; Waaland, J. R., Eds. Algae and Human Affairs; Cambridge University Press: New York, 1988. (4) Horacek, J.; Soukupova, L.; Puncochai, M.; Slezak, J.; Drahos, J.; Yoshida, K.; Tsutsumi, A. J. Hazard. Mater. 1994, 37, 69-76. (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 275-287. (6) Gosset, T.; Trancart, J.-L.; Thevenot, D. R. Water Res. 1986, 22, 21-26. (7) Albasel, N.; Cottenie, A. Soil Sci. Soc. Am. J. 1985, 49, 386-389. (8) Bunzl, K.; Schmidt, W.; Sansoni, B. J. Soil Sci. 1976, 27, 32-41. (9) Sposito, G. Chemical Equilibria and Kinetic in Soils; Oxford University Press: New York, 1994. (10) McBride, M. B. Environmental Chemistry of Soils; Oxford University Press: New York, 1994. (11) Sparks, D. L. Soil Physical Chemistry; CRC Press: Boca Raton, FL, 1986. (12) Morel, F. M. M. Principles of Aquatic Chemistry; WileyInterscience: New York, 1983; pp 266-293. (13) Kadlec, R. H.; Keolian, G. A. In Peat and Water; Fuchsman, C. H., Ed.; Elsevier: London, 1986; pp 61-93. (14) Ref 5, pp 275-287. (15) Crist, R. H.; Martin, J. R.; Carr, D.; Watson, J. R.; Clarke, H. J.; Crist, D. R. Environ. Sci. Technol. 1994, 28, 1859-1866. (16) Bloom, P. R.; McBride, M. B. Soil Sci. Soc. Am. J. 1979, 43, 687692. (17) Andre, J. P.; Pijarowski, L. J. Soil Sci. 1977, 28, 573-584. (18) Gamble, D. S.; Schnitzer, M.; Kerndorff, H. Geochim. Cosmochim. Acta 1983, 47, 1311-1323 and references cited therein. (19) Kunin, R. Ion Exchange Resins, 2nd ed.; R. E. Kreiger Publishing Co.: Malabar, FL, 1958; p 337. (20) Gamble, D. S. Can. J. Soil Sci. 1989, 69, 313-324. (21) Helfferich, F. Ion Exchange; Dover Publications: New York, 1995; pp 151-156. (22) Ref 10, pp 63-117. (23) Marinsky, J. A. In Aquatic Surface Chemistry; Stumm, W., Ed.; Wiley-Interscience: New York, 1987; pp 71-72. (24) Cotton, F. A.; Wilkinson, G. Advanced Inorganic Chemistry, 5th ed.; Wiley-Interscience: New York, 1988; pp 45-47. (25) Frausto da Silva, J. J. R. J. Chem. Educ. 1983, 60, 390-392. (26) Poonia, N. S.; Bajaj, A. V. Chem. Rev. 1979, 79, 410-415. (27) Percival, E. Br. Phycol. J. 1979, 14, 103-117. (28) Lowry, T. H.; Richardson, K. S. Mechanism and Theory in Organic Chemistry, 3rd ed.; Harper and Row: New York, 1987; pp 143158. (29) Crist, R. H. Messiah College, unpublished results.

Received for review July 31, 1995. Revised manuscript received March 8, 1996. Accepted April 29, 1996.X ES950569D X

Abstract published in Advance ACS Abstracts, June 15, 1996.

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