Research Irreversible Adsorption of Naphthalene and Tetrachlorobiphenyl to Lula and Surrogate Sediments AMY T. KAN,* GONGMIN FU, M. A. HUNTER,† AND MASON B. TOMSON Hazardous Substance Research Center/South & Southwest, Department of Environmnetal Science and Engineering, Rice University, MS-317, 6100 S. Main Street, Houston, Texas 77005-1892
Several unique features of sorption irreversibility have been investigated in this paper. Adsorption has been found to be biphasic with about 30-50% of the adsorbed mass residing in the irreversibly sorbed compartment, until this compartment is filled, and the rest of the mass resides in the labile compartment. Naphthalene in the reversible compartment follows a linear adsorption isotherm with a normal organic carbon-based partition coefficient. A finite fixed total compartment size is observed for the irreversible irr (µg/g), on both natural and surrogate solids. fraction, qmax In multiple batch adsorption/desorption experiments, the irr maximum concentrations that resist desorption are qmax ≈ irr 10 µg/g for naphthalene on Lula sediment and qmax ≈ 0.36 µg/g for 2,2′,5,5′-tetrachlorobiphenyl (2,2′,5,5′-CB) on both Lula and surrogate solids. The concentration in the irreversibly sorbed compartment varied with the initial naphthalene concentration available for adsorption. In addition, the amount in the irreversibly sorbed compartment increases linearly with the number of adsorption steps irr is reached. After the until the maximum concentration qmax maximum concentration of the irreversibly sorbed compartment is satisfied, the adsorption/desorption of naphthalene and 2,2′,5,5′-CB becomes reversible. The irreversibly sorbed compartment appears to be at equilibrium with the aqueous phase when the labile naphthalene or 2,2′,5,5′-CB is removed, but the equilibrium concentration is much lower than would be predicted with conventional hydrophobic partitioning theory. The aqueous phase concentration in equilibrium with the irreversibly sorbed compartment is about 2-5 µg/L for naphthalene and 0.05-0.8 µg/L for 2,2′,5,5′-CB. Similar adsorption/desorption phenomena are observed with both a natural sediment and a well-characterized sorrogate solid.
Introduction The distribution of persistent organic compounds in a contaminated site is strongly influenced by the sorptive behavior of chemicals with soil and/or the sediment. In order * Corresponding author phone: (713)285-5224; fax: (713)285-5203; e-mail address:
[email protected]. † Present address: Connecticut Agricultural Experimental Station, Dept. of Soil and Water, P. O. Box 1106, New Haven, CT 06504.
2176
9
ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 31, NO. 8, 1997
to set realistic soil quality limits for regulatory purposes and to estimate and improve the efficiency of remediation techniques, the mechanisms that govern contaminant dissipation from soil and sediments have to be fully understood (1). Extensive studies have been devoted to understanding the thermodynamics and kinetics of the adsorption process (2-13). The results of many laboratory and field observations indicate that a fraction of organic pollutants resists desorption and/or that the desorption rate is very slow (4, 6, 14-22). The desorption of contaminants from sediments and soil is typically biphasic, whereby the majority of contaminants desorbs within a short time frame [1-7 day (4)] and the removal of the remaining fraction requires weeks to months and often longer. Other processes that govern the loss of contaminants from soil (e.g., volatilization, leaching, biodegradation, and chemical degradation) also exhibit similar biphasic patterns (1, 10, 23, 24). Diffusional mass transfer mechanisms, i.e., intraparticle and/or intraorganic matter diffusion, are frequently invoked to be the primary rate-limiting factors (9, 15, 17, 25). However, the diffusional mass transfer mechanisms often fail to make a priori prediction of contaminant release based on measured physical properties of the solids and chemicals. For example, Hutchins et al. (26) observed that the groundwater beneath the infiltration sites contained numerous trace level contaminants that were released at a constant concentration. Steinberg et al. (21) found a soil fumigant (1,2-dibromoethane) in agricultural topsoil 19 years after its last known application, despite its high volatility and degradability. Pereira et al. (27) found that the concentration of halogenated organic compounds in native water, suspended sediment, and biota was far below the values predicted with respect to concentrations in the contaminated bottom sediment collected from Bayou d’Inde, LA. Similarly, McGroddy and Farrington (28) as well as Readman et al. (29) observed a fraction of PAHs in river sediment not available for partitioning. These observations prompt the need for a better understanding of the complete desorptive behavior of organic pollutants in sediment/water systems. A complete description of adsorption and desorption would have to take into account numerous phenomena, including surface reactions, solution phase reactions, and mass transport phenomena. Pignattelo and Xing (30) reviewed recent research into the causes of slow sorption and desorption rates. A short list of properties that might complicate the interpretation of adsorption and desorption includes the following: the presence of different functional groups, hydrogen or electronegative atoms to form hydrogen bonds, surface charge, chemical solubility, soluble complex formation in solution, the presence of steric hindrances from fine pores of varying size, and complex microscopic hydrodynamic conditions near particle surfaces. Numerous recent papers have considered such effects with both new observations and mathmetical models to interpret the results. For example, Carroll et al. (31) studied the desorption of contaminated Hudson River sediment and reported that a significant fraction of the adsorbed mass (45%) resists desorption. Milling the sediment had no effect on 2,2′,5,5′CB desorption, while pretreatment of the sediment using heat enhanced the release of 2,2′,5,5′-CB from the resistant fraction. They attributed the cause of slow release to be the slow desorption from a condensed phase of soil organic matter.
S0013-936X(96)00195-2 CCC: $14.00
1997 American Chemical Society
TABLE 1. Testing Parameters for Naphthalene and PCB Adsorption/Desorption Experiments solution matrix solid/solution ratios mixing desorption exps 1-3 exp 4 no. of adsorption/desorption steps exp 1 exp 2 exp 3 exp 4 naphthalene concentration (C0) used for adsorption exp 1 exp 2 exp 3 exp 4
Naphthalene Experiments 1 mM CaCl2‚2H2O, 0.1 mM MgCl2, 0.5 mM Na2B4O9‚10 H2O (pH 8.0), 0.01 M formaldehyde 2 g of Lula sediments, 24.5 mL of electrolyte solution horizontal shaking 60% supernate was replaced with clean electrolyte 90% supernate was replaced with clean electrolyte 6 adsorption/desorption cycles with 5-7 desorption steps per adsorption 3 adsorption/desorption cycles with 5-13 desorption steps per adsorption 2 adsorption/desorption cycles with 5 desorption steps 1 adsorption/desorption cycle with 20 desorption steps (1) 2.844, (2) 2.357, (3) 3.182, (4) 3.256, (5) 2.963 mg/L in each adsorption cycle (1) 6.307, (2) 6.044, (3) 9.675 mg/L in each adsorption cycle (1) 15.173, (2) 14.255 mg/L in each adsorption cycle (1) 14.197 mg/L PCB Experiments
solution matrix exp 5 exp 6 solid/solution ratios exp 5 exp 6 mixing desorption exp 5 exp 6 no. of adsorption/desorption steps exp 5 exp 6 PCB concentration (C0) used for adsorption exp 5 exp 6
0.1 M NaCl, 0.01 M sodium acetate, 0.01 M sodium azide, pH 2.5 0.15 M NaCl, 0.025 M sodium phosphate, 0.01 M sodium azide, pH 5.5 4 g of surrogate sediments, 40 mL of electrolyte solution 4 g of Lula sediments, 40 mL of electrolyte solution horizontal shaking 90-95% supernate was replaced with clean electrolyte 90-95% supernate was replaced with clean electrolyte 8 adsorption/desorption cycles with 8-15 desorption steps per adsorption 6 adsorption/desorption cycles with 8-39 desorption steps per adsorption (1) 10.59, (2) 10.56, (3) 10.39, (4) 9.89, (5) 11.65, (6) 9.14, (7) 9.38, (8) 11.25 µg/L in each adsorption (1) 12.12, (2) 12.82, (3) 12.90, (4) 12.76, (5) 12.52, (6) 12.91 µg/L in each adsorption
Connaughton et al. (32) proposed a continuum of rate-limiting compartments to model the observed persistant release of naphthalene from soil. Burgos et al. (24) observed about 2-8% of naphthalene trapped within mass transfer-limited sites, but not covalently bound. Karimi-Lotfabad et al. (33) observed the polymerization of anthracene in dry soil. Young and Weber (34) showed that organic matter from shale adsorbed chemicals much stronger than surface soil and suggested that the differences in adsorption were due to differences in microcrystallinity of the organic phase as a result of diagenetic alteration. Weber and Huang (35) proposed that the organic matter on soil had both amorphous and microcrystalline phases and the adsorption of contaminants to the different domains account for the time-dependent adsorption observed in the bulk solid. McGroddy et al. (36) observed the resistant desorption of PAHs from Boston Harbor sediment, while the PCBs desorption from the same sediment were consistent with equilibrium partitioning model. They attributed the observed PAHs desorption resistance to be the result of some PAHs strongly adsorbed to highly aromatic soot particles in sediment. It will be shown below that adsorption and desorption of 2,2′,5,5′-CB and naphthalene cannot be interpreted in the same manner. Sorption irreversibility was previously proposed by the present authors (4, 11) but remains a debatable issue. A minimum thermodynamic requirement for adsorption and desorption to be irreversible is that there be a physicalchemical rearrangement in the solid phase after adsorption occurs, i.e., the desorption takes place from a different molecular environment than that to which it adsorbed (37). The term “irreversible” used herein is in the same manner as it is commonly used in the physical-chemical literature, e.g., Adamson (37), Bailey et al. (38), and Burgess et al. (39).
Specifically, irreversible is meant to imply that when adsorption to sediment OC occurs, the sorption causes a rearrangement of the solid or OC matrix in an irreversible manner and that desorption from this altered solid or OC matrix is not the reverse of the adsorption process. Experimentally, it is difficult to differentiate sorption irreversibility from various other reactions as mentioned above. In this paper, the irreversible adsorption phenomenon is characterized with efforts directed toward the elucidation of the causes of the observed phenomena using multiple cycles of adsorption/desorption. Due to the complex nature of the natural sediment and its organic phase, parallel adsorption/ desorption experiments were conducted with a well-behaved surrogate sediment. The surrogate material is a surfactant hemimicelle-coated nonporous anatase developed by Hunter et al. (40). The testing with surrogate sediment is to eliminate the uncertainty associated with the natural sediment, namely, the heterogeneity and microporosity of the organic matter and the associated inorganic matrix. The combination of multiple applications of contaminants and comparison of the adsorption/desorption of 2,2′,5,5′-CB to both natural and surrogate sediments have provided insightful information regarding the adsorption/desorption process.
Materials and Methods Sorbent. Sediment from Johnson Ranch, Lula, OK, was used in this study. The ranch is located near the margin of the flood plain of a small river. Sediment was air-dried, sieved through cheesecloth to remove vegetative matter and pebbles, and stored in the refrigerator. The sediment contains 92% sand, 6% silt, and 2% clay with particle sizes ranging from 2 mm to less than 0.063 mm and a weighted average diameter
VOL. 31, NO. 8, 1997 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
9
2177
of 0.23 mm. It contains 0.27% sediment organic matter. About half of the organic matter resides with the larger sand grains at a concentration of 0.17-0.29% OC, and the silt and clay fractions have a slightly higher organic carbon content (0.390.75% OC for 98% as determined by GC/MSD. Upon receipt of the radiolabeled chemicals, stock solutions of 130 (for naphthalene with specific activity of 10.1 µCi/µmol) and 53 mg/L (for naphthalene with specific activity of 49.8 µCi/ µmol) [14C]naphthalene solution and 17.4 mg/L [14C]2,2′,5,5′CB solution were prepared in methanol. The [14C]naphthalene was further diluted with unlabeled [12C]naphthalene stock solution (1500 mg/L in methanol) to yield a specific activity of 0.680 (exps 1 and 2), 0.238 (exp 3), and 2.262 (exp 4) µCi/ µmol, respectively. The amount of methanol and formaldehyde added to the vials constitutes less than 2% (v/v) or 1% (mole fraction) of the liquid, which should not affects the properties of the liquid phase significantly. Chemicals used in this research were reagent grade or better, except for sodium dodecyl benzenesulfonate. The commercial anionic surfactant used as the adsorbed organic carbon source for surrogate was sodium dodecyl benzenesulfonate (C12H25C6H4SO3- Na+, Rhodocal DS 10, RhonePoulenc). Formaldehyde (0.01 M) or sodium azide (0.01 M) was added to the solution as a bacterial inhibitor. Naphthalene solutions of 2.39-15.1 mg/L were prepared before each adsorption experiment using the [14C]naphthalene stock solution and the pH-buffered electrolyte solution shown in Table 1. 2,2′,5,5′-CB solutions between 9.14 and 11.65 µg/L were prepared in electrolyte solution using [14C]2,2′,5,5′-CB stock solution and the pH-buffered electrolyte solution shown in Table 1. Batch Adsorption and Desorption Experiments. Table 1 lists the ranges of parameters tested for the adsorption/ desorption experiments. The adsorption/desorption experiments were conducted in a batch reactor, which consisted of a glass vial of total volume approximately 26 mL (naphthalene experiments) or 40 mL (2,2′,5,5′-CB experiment) and sealed with Teflon-faced silicone septum (Wheaton). Control experiments were run with the vials containing naphthalene solution at 1 mg/L to test the integrity of the vial. Since 2,2′,5,5′-CB has the tendency to adsorb to glasswares, glasswares were cleaned by the procedure discribed by Hunter et al. (40). Less than 5% reduction of the solution phase concentration was observed over a 30-day period. At the beginning of the adsorption experiments, dried sediment was added to the vial before the addition of naphthalene solution
2178
9
ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 31, NO. 8, 1997
FIGURE 1. Plot of the solid phase concentration versus solution phase concentration of naphthalene for the first cycle of adsorption/ desorption data from exps 1-4 of Table 1. The four open squares are the initial adsorption points from exps 1-4. The open circles and the line are the adsorption isotherm data reported in Kan et al. (4). at concentrations that ranged from 9 µg/L to 15.1 mg/L. After filling with sediment and solution, the vial typically contained a very small headspace of less than 0.1 mL. The sediment/ water mixture was horizontally mixed in a shaker bath at room temperature. After 24 h, the sediment was separated from solution by centrifugation at a centrifugal force of 300g for 15 min (International Clinical Centrifuge, W. H. Curtin Co.), and the chemical concentration in supernate was analyzed. Numerous other centrifugation rates and separation methods were tested with no appreciable effect (11). After adsorption, desorption was induced by successively replacing 60-95% supernate with naphthalene or 2,2′,5,5′CB-free electrolyte solution. The incubation period for desorption experiments was 1-3 days, with exceptions where incubation periods up to 6 months were used. From 5 to 39 successive desorption steps were employed in these experiments. After the successive desorption steps, the solution was decanted and wet sediment was left in the bottle with about 0.5 mL of liquid remaining (actual amounts were determined by weight). To conduct another cycle of adsorption/desorption experiments, an aliquot of 24.5 mL of naphthalene or 40 mL of 2,2′,5,5′-CB solution was added to the vial. The adsorption/desorption cycles were repeated for 1-8 adsorption/desorption cycles. The solution phase naphthalene and 2,2′,5,5′-CB concentrations were determined by scintillation counting (Beckman LS3801) using Beckman ReadySafe scintillation cocktail. The concentration of chemicals in both the liquid and solid phases are within the sensitivity of scintillation counting. At counting times of 10 min per sample, the 95% confidence interval of the counting efficiency is about (5% at 0.635 µg/L 2,2′,5,5′-CB and (9.8% at 0.027 µg/L 2,2′,5,5′-CB concentrations. The solid phase naphthalene and 2,2′,5,5′-CB concentrations were calculated by assuming that the changes in solution phase naphthalene and 2,2′,5,5′-CB concentrations during adsorption or desorption are equal to the changes in solid phase concentration. This mass balance assumption has been rigorously tested in similar previous research by
FIGURE 2. Plots of the (A) solution phase and (B) solid phase naphthalene concentration profiles versus adsorption/desorption steps for exps 1-4. Desorption time is 1-3 days for most adsorptions/desorptions. Several desorptions were conducted at longer desorption time, and the times (days) were marked individually by the particular desorption step. Kan et al. (4). The solid phase concentrations of naphthalene and 2,2′,5,5′-CB were experimentally confirmed by solvent extraction for exps 1 and 5. The solid phase naphthalene concentration of exp 1 was measured at the end of 33 adsorption/desorption steps by methylene chloride extraction, where solvent was in reflux with a round-bottom flask containing sediment for 24 h in a constant temperature water bath at 45 °C. The extraction efficiency of this procedure has been reported previously (4). The 2,2′,5,5′-CB/surrogate sediment was extracted with acetone in three 24-h successive extraction steps. The concentrations of 2,2′,5,5′-CB and naphthalene recovered from the solvent extraction were quantified by scintillation counting. The identities of the chemicals were verified to be naphthalene and 2,2′,5,5′tetrachlorobiphenyl by GC/MS and GC/ECD analysis, respectively.
Results Multiple Adsorption/Desorption of Naphthalene to Lula Sediment. In this study, an individual experiment consisted of repetitive cycles of adsorption/desorption. A cycle of adsorption/desorption consists of adsorption followed by several successive desorption steps. In Figure 1, the first cycle of naphthalene adsorption/desorption experiments for exps 1-4 is compared to those reported previously (4). The equilibrium adsorption data for exps 1-4 can be fitted to the linear isotherm reported previously [Kp ) 2.2 or Koc ) 102.70 (4)]. Desorption irreversibility, similar to that observed in Kan et al. (4), was again observed regardless of the higher naphthalene concentration. If ideal, linear, and reversible
adsorption/desorption is assumed as a point of reference, over 97% of the naphthalene should have been desorbed within five to six desorption steps for exps 1-4 [see Kan et al. (4)]. Therefore, it is proposed that the naphthalene remaining in the solid phase after five desorption steps is irreversibly sorbed. The mass in the “irreversibly sorbed” fraction is about 50% of the total adsorbed mass. This mass fraction that becomes irreversibly sorbed is similar in range to that reported previously (4, 11). The solution phase and solid phase naphthalene concentration profiles of the four naphthalene experiments at every adsorption/desorption step are plotted in Figure 2. Both the solution phase and solid phase concentrations increase to a peak value during adsorption, and then the concentrations decline in the subsequent successive desorption steps. Extensive desorption times of up to 74 days do not affect the solution phase and solid phase equilibrium concentrations substantially (Figure 2A). The irreversible compartment saturates with naphthalene at about 12 µg/g (Figure 2B, e.g., exp 4, Steps 10-20). The calculated solution phase concentration would be C ) 6 µg/mL at equilibrium with 12 µg/g naphthalene on solid phase (Kp ) KocOC = 2 mL/g). Yet, the corresponding observed solution phase concentration (Figure 2) was near 0.003 µg/mL, or about 2000 times lower than expected. The solution phase concentration does not increase with lengthy desorption time, despite the fact that the solution is in equilibrium with a relatively high solid phase naphthalene concentration. Therefore, the observed phenomenon is probably not a simple kinetic or a diffusion limited process.
VOL. 31, NO. 8, 1997 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
9
2179
FIGURE 3. Plot of the (A) solution phase and (B) solid phase PCB concentration profiles versus adsorption/desorption steps for the surrogate adsorbents. In panel A, the circles are the initial PCB concentration (C0) used in the eight adsorption steps, and the dashed horizontal lines represent the mean concentrations of the PCB solution used for adsorption experiments (C0) and the resulting solution phase concentration (Cads) at the eight adsorption steps. Desorption time is 1-3 days for most adsorption or desorption steps. Several desorptions were conducted at longer desorption time, and the times (days) were marked individually by the particular desorption step. The solid and dotted lines are the average and 95% confidence interval of the residual concentrations observed in each desorption cycles. In panel B, the intervals between the dotted and dashed lines represent the reduction in solid phase concentration due to reversible desorption in each adsorption/desorption cycle. The intervals between the two dashed lines represent the solid phase concentration that become irreversibly sorbed for each cycle. At each adsorption/desorption cycle, up to a fixed limit, a portion of the adsorbed naphthalene resists desorption. Note that the solid phase concentration that resists desorption reaches a constant value of 10-12 µg/g sediment after multiple cycles of adsorption and desorption for these four experiments. These data show that the irreversibly sorbed compartment is finite in size and that it is about 7.8 × 10-8 mol/g sediment (2.9 × 10-5 mol/g organic carbon) or 3.7 mg of naphthalene/g of organic carbon. Once the irreversible capacity is reached, the additional mass of naphthalene adsorbed in subsequent cycles desorbs within five desorption steps as would have been predicted by equilibrium partitioning and assuming a measured Kp = 2 mL/g. For example, in the last cycle of exps 1-3, 97-102% of the added naphthalene desorbed within five desorption steps. The sediment from exp 1 (Step 33 in Figure 2B) was extracted with methylene chloride, and 9.385 µg/g of the irreversibly sorbed naphthalene fraction was recovered in the solvent phase as determined by scintillation counting and capillary GC analysis. The amount recovered in the methylene chloride extract represents a mass balance of 103%. In other words, the naphthalene that disappears from the aqueous phase was quantitatively extracted from the solid phase and verified to be the chemically unaltered naphthalene via GC/MS analysis.
2180
9
ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 31, NO. 8, 1997
Previously, similar recovery efficiencies were found by the authors with single adsorption followed by multiple desorption (4). The data of exp 4 show that naphthalene is released from the irreversibly sorbed compartment at a low and nearly constant solution concentration (see Figure 2A). In exp 4 the initial solution phase concentration was 14.197 mg/L, which from previous experience was expected to saturate the irreversible compartment in one adsorption step. Experiment 4 is similar to exp 3 except that 90% versus 60% the of supernate was replaced in each desorption step in an attempt to accelerate the desorption, and up to 20 desorption steps were used to remove the labile naphthalene from the system. Although naphthalene desorbed slightly faster in exp 4 than in exp 3 for the first three desorptions, only a small mass of naphthalene continues to desorb per desorption step after the third desorption step. After the labile naphthalene was desorbed during the first five desorption steps (Figure 2), the solution phase naphathalene concentration approached a nearly constant concentration of about 2-5 µg/L for the next 10-15 desorption steps. Furthermore, during these latter 10-15 desorption steps, the solution phase concentration does not vary appreciably for incubation times from 3 to 29 days.
FIGURE 4. Plots of the (A) solution phase and (B) solid phase concentration profiles versus adsorption/desorption step number for the six multiple cycles of PCB adsorption/desorption with the Lula sediments. In panel A, the circles are the initial PCB concentration (C0) used in the adsorption steps, and the dashed horizontal lines represent the mean concentrations of the PCB solution used for adsorption experiments (C0) and the resulting solution phase concentration (Cads) at the six adsorption steps. The time (days) is the equilibrium time used in the corresponding experiments. The solid lines are the average residual concentrations observed in each desorption cycle. Irreversible Adsorption/Desorption of 2,2′,5,5′-CB to both Lula and Surrogate Sediments. Hunter et al. (40) have developed a surrogate sediment material (a dodecyl benzene sulfonate-coated anatase) that exhibits similar adsorption/ desorption characteristics as natural sediment. The adsorption of 2,2′,5,5′-CB to surrogate can be approximated by a linear isotherm with an organic carbon-based adsorption coefficient (Koc, the linear isotherm slope) of 104.34 (mL/g of OC). The adsorption data for 2,2′,5,5′-CB to Lula sediment are closely related to the linear isotherm determined with surrogate solid (40). Interestingly, the desorption of 2,2′,5,5′CB from surrogate also exhibits a similarly strong irreversible desorption as that of Lula, indicating that the observed irreversible adsorption phenomenon is not a unique phenomenon with the natural sediment. Figure 3 is a plot of the solution phase and solid phase 2,2′,5,5′-CB concentrations versus step number for the multiple adsorption/desorption surrogate experiment (exp 5). Eight adsorption/desorption cycles were conducted. The observed solution phase concentration reduces exponentially in the first three to five desorption steps and then levels off to a constant sub-ppb concentration in the following desorption steps. A significant fraction of 2,2′,5,5′-CB resists desorption. In Figure 3A, the circles are the 2,2′,5,5′-CB concentrations (C0) used in the eight adsorption experiments.
Note that solution phase 2,2′,5,5′-CB concentrations (Cads) reaches an average concentration of 1.78 ( 0.35 µg/L for the eight adsorption steps. Therefore, nearly equal amounts of 2,2′,5,5′-CB are adsorbed in each adsorption step, even though 2,2′,5,5′-CB from previous adsorption steps remains. This observation suggests that the 2,2′,5,5′-CB in the irreversible compartment does not interact with the solution phase during the subsequent adsorption steps. It is noteworthy that the solution phase concentration tends to level off at a constant sub-ppb concentration of 0.050.8 µg/L after a few successive desorptions, phenomenologically similar to that observed in the naphthalene experiment (exp 4). The solid and dotted horizontal lines in Figure 3A are the average and 95% confidence intervals of the residual solution phase concentration for the last four to nine desorption steps following an adsorption cycle. The residual solution phase concentration reaches a constant following each adsorption cycle, but this concentration increases with the subsequent adsorption cycles. Extended desorption times of between 9 and 27 days do not change the solution phase concentration substantially. To confirm that the solid phase 2,2′,5,5′-CB concentration was correct and that 2,2′,5,5′-CB was not altered chemically due to photocatalytic degradation by anatase, the sample was extracted with acetone at the end of the experiment (step
VOL. 31, NO. 8, 1997 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
9
2181
FIGURE 5. Schematic representation of the proposed biphasic irreversibly adsorption model. 85). The extract was analyzed on capillary GC/ECD, and the GC chromatogram produced a peak with retention time identical to that of [12C]-2,2′,5,5′-tetrachlorobiphenyl standard solution. The amount of 2,2′,5,5′-CB recovered in the acetone extract was 82% of the calculated solid phase concentration. Note that the mass balance of this experiment is not as good as that for the naphthalene experiments. However, this experiment is much more demanding than the naphthalene experiments. The lower mass balance is probably caused by a combination of lower 2,2′,5,5′-CB concentration and more desorption experiments; the experiment consisted of the replacement of 85 aliquots over about 6 months. The significance of the finding is that the majority of the chemical that disappeared from the aqueous phase was still present in the solid phase as the unaltered 2,2′5,5′-tetrachlorobiphenyl at the end of 85 adsorption/desorption steps and several months of experimentation. As shown in Figure 3, a significant fraction of the adsorbed 2,2′,5,5′-CB resists desorption in the first six adsorption cycles, and the concentration of 2,2′,5,5′-CB in the irreversible compartment increases stepwise, in a manner similar to naphthalene adsorption/desorption. In the last two cycles, the amount of 2,2′,5,5′-CB desorbed in nine steps is similar to the amount adsorbed. Therefore, the irreversible compartment is essentially saturated in six to seven adsorptions, and subsequent adsorption/desorption cycles are reversible, similar to that observed in naphthalene experiments (exps 1-3). The maximum solid phase 2,2′,5,5′-CB concentration that resists desorption is 0.35 µg/g. In Figure 4 are plots of solution phase and solid phase 2,2′,5,5′-CB concentration profiles for the corresponding Lula sediment experiments. The organic carbon content of the Lula sediment is slightly lower than that of the surrogate material, and the 2,2′,5,5′-CB concentration used for the experiment is slightly higher than that used in the surrogate experiment. The observed solution phase and solid phase concentration profiles have similar characteristics as the surrogate solid. The solution phase concentration (Cads ) 1.85 ( 0.24 µg/L) remains nearly constant over the six adsorption steps. During desorption, the solution phase concentration decreased exponentially and reached a residual solution phase concentration of 0.2-0.6 µg/L. This residual concentration increases in the subsequent adsorption/ desorption cycles. According to the surrogate experiment, the irreversible compartment is filled in about six adsorption cycles. When examining the sixth cycle adsorption/desorption data, 90% of the adsorbed mass desorbed in nine
2182
9
ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 31, NO. 8, 1997
FIGURE 6. Plot of the observed solid phase naphthalene concentration (qtotal) versus the solution phase naphthalene concentration (C) for the five multiple cycles of adsorption/desorption of exp 1. The dashed lines are linear fit of the first five desorptions from each cycle. The value at the Y intercept is assumed to be equal to the solid phase concentration of the irreversibly sorbed compartment (qirr). desorption steps, which approached complete desorption. Note that one desorption step was conducted over 180 days and that the solution phase concentration did not change from that of the previous desorption. The stepwise increase in the concentration of the irreversible compartment is similar to that which occurred with the surrogate and with naphthalene experiments (Figure 4B). The maximum solid phase 2,2′,5,5′-CB concentration that resists desorption is 0.36 µg/ g. Desorption from Lula sediment was carefully examined with 39 desorption steps following the sixth adsorption. The solution phase concentration reached a reasonably constant concentration following six to eight desorption steps. However, the constant solution phase concentration reduced stepwise every few desorption steps, almost mimicking the change in residual concentration during the cyclic adsorption process, but in reverse order. The changes in solution phase concentration after multiple desorptions are reflected in the changes in solid phase concentrations (Figure 4B). The slope of the decline reduces from about 8.2 ng/g per desorption
FIGURE 7. Plot of the solid phase naphthalene concentration of the “labile” compartment (qrev) versus the solution phase naphthalene concentration (C) of all adsorption/desorption data of exps 1-4 and the data from exps 1d, 16d, and 17d of Kan et al. (4). step at the first seven desorptions to about 2.7 ng/g per desorption step at the last seven desorptions.
Discussion The most significant observations of this study are as follows: (1) the well-defined surrogate adsorbent exhibits similar desorptive characteristics as natural adsorbent; (2) both natural and surrogate solids exhibit a fixed maximum capacity of the irreversibly sorbed fraction; and (3) the desorbed concentrtion is not time dependent over a period of 6 months. No existing mechanism(s) can be used to explain the observations of this study. First, extended desorption time does not change the solution phase concentration substantially, indicating that the observed phenomenon is probably not a simple kinetic or diffusion-limited process. Second, the diffusion-limited transport through either the organic matter or the microporosity of the solid is often proposed to be the mechanism causing low-level residual desorption (17, 30, 31). However, resistant desorption is observed in the surrogate material, which is nonporous and homogeneous in organic composition. Therefore, the observation cannot be due simply to the heterogeneity of organic phases (34-36, 43), nor the diffusion-limited transport (17, 30, 31) mechanisms proposed by others. Thirdly, no existing mechanism can be used to explain the additive nature of the filling of the irreversible adsorption compartment. The observed stepwise increase in the irreversibly sorbed mass has not been reported previously. The stepwise increase in solid phase concentration is similar to the self-assembly of organic substances to a glass slide typically observed in Langmuir-Blodgett-type phenomenon. In order to form multiple layers of LangmuirBlodgett film, multiple dipping of the glass slide to the organic solution is required, and only a monolayer is formed at each dipping (44). Unfortunately, no mechanistic explanation has been forward for the Langmuir-Blodgett phenomenon (45). Based on the data presented in this paper, a conceptual biphasic irreversible adsorption model is proposed (illustrated in Figure 5 for naphthalene). As shown in the schematic representation of the irreversible adsorption process (Figure 5), naphthalene adsorbs to the soil organic matter following the hydrophobic interaction. After adsorption, a portion of
the adsorbed mass becomes irreversibly bound. The phenomenon may be due to the occlusion of chemicals by a cooperative conformational changes of the organic matter during the adsorption process or due to physical rearrangement of the organic matter phase. The conformational or physical rearrangement of the soil in the presence of adsorbed chemicals could cause the chemical environment of the adsorbate to be different and hence be the source of the irreversibility. A portion of the adsorbed mass resides in the irreversible compartment and becomes unavailable for desorption (eq 1), while the adsorbed mass in the “labile” compartment follows the commonly invoked adsorption model (eq 2)
qrev ) qtotal - qirr
(1)
qrev ) KpC
(2)
where qtotal is the total solid phase naphthalene concentration (µg/g), qrev is the solid phase naphthalene concentration (µg/ g) in the “labile” compartment, qirr is the solid phase naphthalene concentration (µg/g) in the irreversibly sorbed compartment, Kp is the partitioning constant (mL/g), and C (µg/mL) is the equilibrium solution phase naphthalene concentration. To illustrate how to obtain the value of qirr and qrev, Figure 6 is a plot of adsorption/desorption isotherms of exp 1 with five cycles of adsorption/desorption. In order to test this conceptual model, the value qirr may be linearly extrapolated to the Y axis from the desorption isotherm using the first five successive desorption data (dashed line). The Y intercept is the solid phase concentration of the irreversibly sorbed compartment (qirr) and qrev can be obtained from the difference of qtotal and qirr. Figure 7 is a plot of qrev versus C for all data of exps 1-4 as well as data from Kan et al. (4). Note that the desorption data from multiple cycle experiments with concentrations ranging from 0.1 to 15 mg/L can be fitted with a linear isotherm and can yield a slope of 1.35 mL/g and a correlation coefficient of 0.965. The slope corresponds to an organic carbon-based partition coefficient (Koc) of 102.70. This Koc value is reasonable, comparing to that reported in the literature [Koc ) 102.65-3.17 (46, 47)].
VOL. 31, NO. 8, 1997 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
9
2183
FIGURE 8. Plot of the solid phase naphthalene concentration of the irreversible fraction versus the number of the adsorption cycle for exps 1-3 of this paper and exps 1d, 16d, and 17d of Kan et al. (4). Although a linear isotherm may be an oversimplification of the adsorption process, the approach appears to satisfactorily repesent the partitioning of naphthalene in the reversible compartment for over 3 orders of magnitude in concentration. However, an appropriate model to characterize the irreversible compartment is less obvious. Figure 8 is a plot of qirr versus the number of adsorption steps. Three data points from a previous paper [exps 1d, 16d, 17d (4)] were included for comparison. The naphthalene concentration used in these adsorption experiments (C0) are 1.14, 0.24, and 0.55 µg/ml, respectively. The increase in solid phase concentration in the irreversibly sorbed compartment after a single adsorption cycle (∆qirr/cycle, the slope of Figure 8) appears to be proportional to the free naphthalene concentration (C0) used in the adsorption experiment. The fraction of adsorbed naphthalene that becomes irreversibly sorbed is about constant for the multiple adsorption steps before the capacity of the irreversibly sorbed compartment is reached. Note that only about 30-50% of the adsorbed mass can be irreversibly bound in any given adsorption step and any concentration (C0). Only when the initial concentration (C0) is sufficiently high, the capacity of the irreversible compartment can be filled in one step. For naphthalene, Lula sediment is saturated with naphthalene at C0 ) 15 mg/L in one desorption step. The observation is consistent with the conceptual model proposed in Figure 5 that a fraction of the adsorbed naphthalene becomes occluded or is converted to some sort of inert form on the solid phase during the adsorption process. As shown in Figure 8, the irreversible compartment has a finite capacity (qirr max) ranging from 10 to 12 µg/g for naphthalene. This finite-limited compartment may be a characteristics of both the solute and the soil organic matter. After the irreversible compartment has been filled, subsequent adsorption steps appear to desorb completely and with a reasonable Kp value that is characteristic of normally reported Kow/Koc relationships. It is proposed that the irreversibly adsorbed naphthalene is in an environment that is energetically more stable than that normally associated with the typical Kow/Koc-type hydrophobic effect-related adsorption. Using the data of exp 4, the magnitude of the increased stability energy of naphthalene can be estimated via eq 3:
µg/mL ) -18 kJ mol (0.003 6.0 µg/mL )
∆G h ° ≈ RT ln
2184
9
-1
(3)
ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 31, NO. 8, 1997
Similarly, the magnitude of the increased stability energy for 2,2′,5,5′-CB in surrogate or Lula solid is estimated to be about -6 kJ mol-1. Although a free energy of -6 to -18 kJ mol-1 alters the equilibrium considerably, it does not imply any change in adsorption bonding, such as covalent bond formation, but is consistent with a normal physical-chemical adsorption process (37). The proposed conceptual model of irreversible adsorption is consistent with many laboratory and field observations mentioned above (18, 21, 27-29, 31, 48), even though different mechanisms have been proposed in the respective papers. The data and concept presented in this paper are uniquely different and require new interpretation. The conceptual model proposed in this paper appears to be promising. However, we believe that current level of understanding with respect to the formation of the irreversible fraction is insufficient and should be a direction of future research.
Acknowledgments This research has been conducted under the auspices of the South & Southwest Hazardous Substance Research Center with funding provided by the Office of Exploratory Research of the U.S. Environmental Protection Agency.
Literature Cited (1) Beck, A. J.; Wilson, S. C.; Alcock, R. E.; Jones, K. C. Crit. Rev. Environ. Sci. Technol. 1995, 25, 1-43. (2) Kan, A. T.; Tomson, M. B. Environ. Toxicol. Chem. 1990, 9, 253263. (3) Kan, A. T.; Tomson, M. B. J. Contam. Hydrol. 1990, 5, 235-251. (4) Kan, A. T.; Fu, G.; Tomson, M. B. Environ. Sci. Technol. 1994, 28, 859-867. (5) Brusseau, M. L.; Rao, P. S. C. Crit. Rev. Environ. Control 1989, 19, 33-99. (6) Ball, W. P.; Roberts, P. V. Environ. Sci. Technol. 1991, 25, 12371249. (7) Weber, W. J., Jr.; Miller, C. T. Water Res. 1988, 22, 457-464. (8) Miller, C. T.; Weber, W. J., Jr. Water Res. 1988, 22, 465-474. (9) Karickhoff, S. W. J. Hydraul. Eng. 1984, 110, 707-735. (10) Pignatello, J. J. In Organic Substances in Soils and Water; Beck, A. J., Jones, K. C., Hayes, M. H. B., Mingelgrin, U., Eds.; Royal Society of Chemistry: Cambridge, U.K., 1993; p 128. (11) Fu, G.; Kan, A. T.; Tomson, M. B. Environ. Chem. Toxicol. 1994, 13, 1559-1567. (12) Harmon, T. C.; Roberts, P. V. Environ. Sci. Technol. 1994, 28, 1650-1660. (13) Farrell, J.; Reinhard, M. Environ. Sci. Technol. 1994, 28, 53-62.
(14) Brusseau, M. L.; Jessup, R. E.; Rao, P. S. Environ. Sci. Technol. 1990, 24, 727-735. (15) Ball, W. P.; Roberts, P. V. Environ. Sci. Technol. 1991, 25, 12231236. (16) Gschwend, P. M.; Wu, S. Environ. Sci. Technol. 1985, 19, 90-96. (17) Wu, S.-C.; Gschwend, P. M. Environ. Sci. Technol. 1986, 20, 717725. (18) Pignatello, J. J. Environ. Toxicol. Chem. 1989, 9, 1107 -1115. (19) Pignatello, J. J.; Frink, C. R.; Marin, P. A.; Droste, E. X. J. Contam. Hydrol. 1990, 5, 195-214. (20) Pignatello, J. J.; Ferrandino, F. J.; Huang, L. Q. Environ. Sci. Technol. 1993, 27, 1563-1571. (21) Steinberg, S. M.; Pignatello, J. J.; Sawhney, B. L. Environ. Sci. Technol. 1987, 21, 1201-1208. (22) Pavlostathis, S. G.; Jaglal, K. Environ. Sci. Technol. 1991, 25, 274279. (23) Hatzinger, P. B.; Alexander, M. Environ. Sci. Technol. 1995, 29, 537-545. (24) Burgos, W. D.; Novak, J. T.; Berry, D. F. Environ. Sci. Technol. 1996, 30, 1205. (25) Brusseau, M. L.; Jessup, R. E.; Rao, P. S. C. Water Resour. Res. 1989, 25, 1971-1988. (26) Hutchins, S. R.; Tomson, M. B.; Ward, C. H. Environ. Toxicol. Chem. 1983, 2, 195-216. (27) Pereira, W. E.; Rostad, C. E.; Chiou, C. T.; Brinton, T. I.; Barbar, L. B., II. Environ. Sci. Technol. 1988, 22, 772-778. (28) McGroddy, S. E.; Farrington, J. W. Environ. Sci. Technol. 1995, 29, 1542-1550. (29) Readman, J. W.; Mantoura, R. F. C. Sci. Total Environ. 1987, 66, 73 -94. (30) Pignatello, J. J.; Xing, B. Environ. Sci. Technol. 1996, 30, 1-10. (31) Carroll, K. M.; Harkness, M. R.; Bracco, A. A.; Balcarcel, R. R. Environ. Sci. Technol. 1994, 28, 253-258. (32) Connaughton, D. F.; Stedlinger, J. R.; Lion, L. W.; Shuler, M. L. Environ. Sci. Technol. 1993, 30, 145-1151. (33) Karimi-Lotfabad, S.; Pickard, M. A.; Gray, M. R. Environ. Sci. Technol. 1996, 30, 1145-1151.
(34) Young, T. M.; Weber, W. J., Jr. Environ. Sci. Technol. 1995, 29, 92-97. (35) Weber, W. J., Jr.; Huang, W. Environ. Sci. Technol. 1996, 30, 881. (36) McGroddy, S. E.; Farrington, J. W.; Gschwend, P. M. Environ. Sci. Technol. 1996, 30, 172-177. (37) Adamson, A. W. Physical Chemistry of Surfaces, 5th ed.; John Wiley & Sons: New York, 1990; Chapter 11. (38) Bailey, A.; Cadenhead, D. A.; Davies, D. H.; Everett, D. H.; Miles, A. J. Trans. Faraday Soc. 1971, 67, 231. (39) Burgess, C. G. V.; Everett, D. H.; Nuttall, S. Pure Appl. Chem. 1989, 61, 1845-1852. (40) Hunter, M. A.; Kan, A. T.; Tomson, M. B. Environ. Sci. Technol. 1996, 30, 2278-2285. (41) Wieland, W.; Wehrli, B.; Stumm, W. Geochim. Cosmochim. Acta 1988, 52, 1969-1981. (42) Somasundaran, P.; Fuerstenau, D. W. J. Phys. Chem. 1966, 70, 90. (43) Maruya, K. A.; Risebrough, R. W.; Horne, A. J. Environ. Sci. Technol. 1996, 30, 2942-2947. (44) Blodgett, K. B. J. Am. Chem. Soc. 1935, 57, 1007. (45) Giles, C. H.; Forrester, S. D.; Roberts, G. G. In Langmuir-Blodgett Films; Roberts, G. G., Ed.; Plenum Press: New York, 1990; Chapter 1. (46) Abdul, A. S.; Gibson, T. L. Hazard. Waste Hazard. Mater. 1986, 3, 125-137. (47) Karickhoff, S. M.; Brown, D. S.; Scott, T. A. Water Res. 1979, 13, 241-248. (48) Witkowski, P. J.; Jaffe, P. R.; Ferrara, R. A. J. Contam. Hydrol. 1988, 2, 249-269.
Received for review February 28, 1996. Revised manuscript received March 11, 1997. Accepted March 25, 1997.X ES9601954 X
Abstract published in Advance ACS Abstracts, June 1, 1997.
VOL. 31, NO. 8, 1997 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
9
2185
