Environ. Sci. Technol. 1999, 33, 3110-3118
Sorption of Linear Alkylbenzenesulfonates on Sediment Materials JOHN C. WESTALL,* HUA CHEN, WANJIA ZHANG,† AND BRUCE J. BROWNAWELL‡ Department of Chemistry, 153 Gilbert Hall, Oregon State University, Corvallis, Oregon 97331-4003
Linear alkylbenzenesulfonates (LAS) are anionic surfactants that are used in large quantities in industrial and consumer products. They enter the environment primarily through wastewater and sludge. In this study, the sorption of LAS to the surfaces of sediment particles was investigated as a function of LAS homologue, H+ concentration in solution, Ca2+ concentration in solution, sediment properties, and concentration of solids. Evidence for both hydrophobic and specific or electrostatic interactions was seen. Isotherms were generally nonlinear and were represented well by the Freundlich and the virial (electrostatic) equations. Comparisons of apparent distribution ratios, D (L/kg), for linear portions of the isotherms showed ∆ log D/∆ log nCH2 ≈ 0.4, ∆ log D/∆ log [H+] ≈ 0.17, and ∆ log D/∆ log [Ca2+] ≈ 0.23. The value of D for different sediments seemed to correlate most closely with the organic carbon content (foc) of the sediments. The value of the distribution ratio increased with the concentration of solids in the system; this effect could be explained partially by the concomitant increase in Ca2+ concentration in solution.
Lewis and Suprenant (6) have reviewed the acute toxicities of surfactants to aquatic invertebrates. Toxicities vary widely, but as an illustration of the magnitude, they report 48-h LC50 values of LAS in the range of 1.8-5.6 mg/L for D. Magna, one of the more sensitive aquatic invertebrates. Ionic organic compounds such as LAS have specific chemical and electrostatic interactions with sorbents and other solutes that are not observed for neutral, non-hydrogenbonding, hydrophobic compounds. Thus, the simple hydrophobic partitioning paradigm (7-10) is not directly applicable for these ionic organic compounds such as LAS. First, specific chemical interactions of the ionic functional group with the sorbent often result in nonlinear isotherms, making it more difficult to assess the tendency to sorb on the basis of a single value of log K. Second, interactions of the ionic functional group in solution and on the sorbent make it necessary to consider speciation of the target compound. Thus, solution chemistry (pH, dissolved salts, etc.) may influence the extent of sorption. Third, the electrostatic interactions of solutes with environmental surfaces, which are generally negatively charged, alter the tendency for sorption, favoring cationic > neutral > anionic, for compounds of similar hydrophobicities (11). All of these factors make it difficult to predict the sorption of LAS to environmental surfaces with a model as simple as hydrophobic partitioning. In this investigation, we attempt to elucidate the factors of solution chemistry and sorbent composition that affect LAS distribution between water and sediments. We have examined the sorption of LAS homologues on five characterized sediments, as a function of the properties of the surfactant, the characteristics of the sediments, and the compositions of the solutions. Pure compounds (as opposed to the mixtures of isomers and homologues in commercial products) were used in this study; the homologues were labeled with 14C to facilitate experimentation at low (submicromolar) concentrations and surface coverages.
Introduction
Models
Linear alkylbenzenesulfonates (LAS) are anionic surfactants that are used in large quantities in industrial and consumer products. They enter the environment primarily through wastewater and sludge. Even though these compounds are reasonably degradable and not particularly toxic (e.g., refs 1 and 2), the magnitude of the quantities in use requires that the possible effects of these compounds in the environment be investigated thoroughly. The fate of LAS in wastewater treatment and in receiving waters has been reviewed recently (3). In general, LAS are not particularly strongly sorbed but may nevertheless accumulate on the surfaces of solids in environments with high amounts of suspended solids, such as in wastewater treatment plants. LAS have been found to accumulate in sewage sludge, particularly sludge that has been treated anaerobically (4, 5). LAS may be found in sediments as well, particularly at wastewater discharge or sludge disposal sites. Kimerle (2) has reported on the aquatic and terrestrial ecotoxicology of LAS. In interpretation of the toxicity data, careful attention must be given to the state of the LAS (free or bound to particulate matter) and the route of uptake.
The two isotherms used to describe the data in this study are the Freundlich isotherm and the virial isotherm. The Freundlich equation is
* Corresponding author. Phone: (541) 737-2591; fax: (541) 7372062; e-mail:
[email protected]. † Deceased. ‡ Present address: Marine Sciences Research Center, SUNY, Stony Brook, NY 11794-5000. 3110
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CLAS(w)nK ) CLAS(s)
(1)
where CLAS(w) is the concentration of LAS in water (mol L-1), CLAS(s) is the “concentration” of LAS on the sorbent (mol kg-1), and K and n are constants. The Freundlich equation describes almost all of our data reasonably well. In principle, it has a mechanistic basis as a summation of Langmuir isotherms with a distribution of ln K values (12); however, inversion of the Freundlich isotherm to the underlying distribution of ln K is of no apparent practical value. In practice, there is often covariance between experimentally determined values of K and n, and differences in energies of sorption for different materials may not readily be deduced from the Freundlich K values. Also, the limit of CLAS(s)/CLAS(w), as CLAS(w) approaches 0, is undefined for n < 1. Thus, the Freundlich isotherm is convenient for empirical representation of the data but of little use for questions about mechanisms. The virial equation (13) is similar to a linear isotherm with an exponential factor, which could be interpreted either as a correction for heterogeneity of the surface or as an 10.1021/es9804316 CCC: $18.00
1999 American Chemical Society Published on Web 07/31/1999
TABLE 1. Electrostatic Model for Adsorptiona mass action eq for adsorption of charged species Ci(w) exp[-zFψ/RT]K ) Ci(s) interfacial capacitance relates potential to charge C ) σ/ψ Faraday const relates charge to surface concn σ ) FCi(s)/s virial eq (eq 2) is identical to (T-1) with b ) F2/RTCs Boltzmann factor for electrostatic energy ∆ log K ) [-ziFψ/RT ln 10]
C ) interfacial capacitance Ci(s) ) concn of ion i on the surface Ci(w) ) concn of ion i in solution F ) Faraday const K ) adsorption const R ) ideal gas constant s ) specific surface area T ) temperature zi ) charge on ion i (as multiple of electronic charge) σ ) surface charge ψ ) surface potential (interfacial potential difference)
(T-1) (T-2) (T-3) (T-4) (T-5) F m-2 mol kg-1 mol m-3 C mol-1 m3 kg-1 J mol-1 K-1 m2 kg-1 K C m-2 V
a All quantities are specified here in SI units; other units are sometimes defined and used in the text.
adjustment for the electrostatic energy of adsorption:
CLAS(w) exp[-bCLAS(s)]K ) CLAS(s)
(2)
where b and K are the adjustable parameters. The electrostatic model is formally equivalent to a Helmholtz or constant capacitance model of the electric double layer (14), as illustrated in Table 1. Thus, the virial equation can be interpreted on a mechanistic basis, the ratio CLAS(s)/CLAS(w) does have a finite value as CLAS(w) approaches zero, and values of log K can be interpreted as limiting “intrinsic” values, characteristic of the sorbent-solute interaction at infinitely low sorbate concentrations. Furthermore, the isotherms in this study could be represented very well by this equation.
Methods Materials. The compounds used in this study were 4-(1methylnonyl)benzenesulfonate, 4-(1-methylundecyl)benzenesulfonate, and 4-(1-methyltridecyl)benzenesulfonate, abbreviated here as C-10, C-12, and C-14 LAS. The compounds were uniformly ring labeled with 14C. They were originally synthesized by New England Nuclear and obtained through the Soap and Detergent Association from Procter and Gamble. The specific activities of C-10, C-12, and C-14 LAS were 1.46, 13.4, and 12.9 Ci/mol, respectively. The radiochemical purity of these compounds was greater than 99%, as determined by HPLC. The purity was routinely verified during the course
of the study. The molar masses (molecular weights) of the LAS ions are 297, 325, and 353 g/mol. The sediments used in this study exhibited a relatively wide range of physical and chemical properties, as shown in Table 2. They were obtained from Prof. John J. Hassett of the University of Illinois. Procedures for characterization of the sediments are discussed by Hassett et al. (15). The deionized water used in all experiments was from a Millipore Milli-Q system, with specific resistivity approaching 18 MΩ cm. Equipment. For equilibration and phase separation, a thermostated shaker (New Brunswick Model G-24) and thermostated centrifuge (IEC Model B-20A) were used. Sediment samples were oxidized and labeled compounds converted to 14CO2 with a Harvey Model 300 oxidizer. Radioactivity was determined with a Beckman Model LS 7800 scintillation counter. Determination of Distribution Ratios. Sorption of LAS on sediments was studied through equilibration of LAS solutions with sediments in batch experiments. Before sediments were equilibrated with solutions of LAS, they were washed eight times to reduce the amount of material, such as dissolved or colloidal organic matter, that tends to remain suspended in solution after centrifugation. Association of organic compounds with this nonseparable material (NSM) can affect the observed distributions (16). The sediments were washed by placing 0.050-1.0 g of sediment in a 35-mL Corex glass centrifuge tube, adding 20 mL of deionized water, shaking for 1 h at 500 rpm and 25 °C, and centrifuging for 10 min at 10 000 rpm (11 000 G) and 25 °C. Supernatant was removed by aspiration through a glass pipet, and the procedure was repeated for a total of eight washes, at which point the moist sediments were freezedried, reweighed, and used in subsequent experiments on the distribution of LAS. The freeze-drying step allows a more accurate determination of the mass of washed sediments and was found not to affect the distribution experiments. After the sediments were washed, freeze-dried, and reweighed, they were equilibrated with LAS solutions according to the following procedure. A 20-mL aliquot of electrolyte solution (0.01 M NaN3 or in some cases 0.01 M NaCl) was added to the washed and dried sediments. Before the solutions were spiked with LAS, the tubes were centrifuged at 11 000 G for 10 min; the purpose of this step was to reduce contact of sediments with localized high concentrations during the spiking of LAS. A 10-200-µL spike of LAS in 95% ethanol was added with a prerinsed Hamilton syringe. No significant effect of the volume of ethanol in the spike on distribution was found. The contents of the tubes, which were sealed with PTFE-lined caps, were mixed on the shaker at a speed of 500 rpm at 25 °C, normally for 4 h, after which aqueous and sediment phases were separated by centrifugation at 11 000 G at 25 °C for 1 h. After centrifugation, the concentration of labeled compound was determined in solution, on the sediment, and adsorbed to the wall of the tube. Aqueous phase samples
TABLE 2. Properties of Sedimentsa sediment
org C, %
sand, %
silt, %
clay, %
CEC, mmol/g
pHb (1:1)
pHc (1:20)
surface area,d m2/g
clay mineralogye
EPA-16 EPA-13 EPA-12 EPA-25 EPA-1
1.20 3.04 2.33 0.76 0.22
0.5 20.3 0 41.9 93.9
60.5 27.1 64.6 37.6 0.0
39.0 52.6 35.4 20.5 6.1
0.110 0.119 0.135 0.089 0.011
6.50 6.90 7.63 7.65 7.30
6.76 7.08 7.52 7.57 -
18 13 12 8 0.8
kao, ver, ill kao, ill kao, ill kao, sm, ill
DCB extractable,f µmol/g Fe Al 286 351 311 123 42
73 70 74 25 2
a Reported by Hassett et al. (15). b Solution was 0.01 M NaCl, 1 g of sediment to 1 mL of solution. c Solution was 0.01 M NaN , 1 g of sediment 3 to 20 mL of solution. d Determined courtesy of Jeff Fahey, Teledyne Wah-Chang, Albany, OR. e Major components as determined by X-ray diffraction: ill ) illite; kao ) kaolinite; sm ) smectite; ver ) vermiculite (31). f DCB (dithionite-citrate-carbonate) procedure of Jackson et al. (41) as modified for extraction twice at 25 °C for 16 h instead of 80 °C for 15 min.
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were collected immediately after centrifugation; 0.2-1.0mL aliquots of the liquid were withdrawn with a graduated glass pipet and transferred to scintillation vials containing 10 mL of Beckman HP/b scintillation fluid. The concentration of LAS sorbed to sediments was determined by oxidizing two subsamples (10-200 mg) from each tube at 900 °C in an oxygen stream in the oxidizer. The 14CO2 was trapped directly in a scintillation vial in a mixture of 5 mL of Packard Carbosorb and 5 mL of Permafluor V scintillation liquid. The walls of the centrifuge tube were then brushed clean, 20 mL of 95% ethanol was added to extract LAS adsorbed to the walls, and 1-mL aliquots were sampled as described for the aqueous phase. The activity of the 14C-labeled LAS in each phase was determined by counting at least 10 000 counts and converting to the decay rate through a quench correction (17). Calibrations were performed for each experiment by adding various amounts LAS to each phase and determining the activity. The calibrations were linear, and the counting efficiency was regularly 93%. The overall relative standard deviation for aqueous samples was estimated to be about 2%, and that for sediment samples was about 5%. For every experiment, the aqueous phase, the sediment, and the walls of the containers were analyzed for LAS. The recoveries (amount of 14C found through analysis divided by the amount of 14C added) regularly ranged from 95 to 101%, with a few exceptions, and averaged 98.8% in 30 experiments conducted on different dates. The fraction of LAS on the walls ranged between 1 and 10%; this fraction was lower when greater amounts of LAS were in the system or when greater amounts of LAS were associated with the sediments. The distribution ratio of LAS between the sediment and aqueous phase, D (L/kg), was calculated directly from the experimentally determined sediment and aqueous phase concentrations CLAS(s) (mol/kg) and CLAS(w) (mol/L). Effects of Solution Composition. Experiments were conducted to study the effects of concentrations of H+ and Ca2+ on the sorption of LAS on EPA-12 sediment. For the experiments with H+, various volumes of 1 M NaOH or HCl were added to slurries of sediments in 0.01 M NaN3. The pH was determined before addition of LAS and immediately after aliquots were removed for aqueous phase LAS determination, to quantify the change in pH over the course of the ≈4-h equilibration step. An Orion Model 8102 Ross combination glass electrode, with 4 M KCl outer filling solution, was used for pH determination. The electrode was calibrated against NIST buffers. For the experiments on Ca2+, CaCl2 was added to the 0.01 M NaN3 solution to yield final concentrations of Ca2+ between 0.10 and 1.0 mM. The background concentrations of Ca(II) and other elements that dissolve from the sediments during equilibration were analyzed by inductively coupled plasma emission spectrometry on a Jarrel-Ash ICAP-9000 spectrometer. Sorption and Desorption Kinetics. The sorption and desorption kinetics of C-12 LAS on sediment EPA-12 were found to be reasonably rapid and reversible, requiring less than 4 h to reach apparent equilibrium in either direction. The value of D obtained for the desorption step was within experimental error of the values obtained for the sorption step, when allowance was made for the dependence of D on the concentration of LAS (i.e., the nonlinearity of the isotherm). These results are in agreement with those of Inoue et al. (18), Matthijs and De Henau (19), and Hand and Williams (20). Degradation of LAS. The results from some of the desorption experiments conducted in 0.01 M NaCl solution suggested that 14C-labeled LAS was undergoing transformation to a more soluble, less readily sorbed form. The value of the distribution ratio of LAS between sediment and water 3112
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FIGURE 1. Sorption isotherms of C-10 LAS, C-12 LAS, and C-14 LAS on sediment EPA-12 (at 0.048 kg/L for C-10 and C-14 and 0.014 kg/L for C-12) in 0.01 M NaN3. The lines were calculated from the virial equation (eq 2) with the parameters in Table 3. approached zero with time, and up to 1% of the 14C could be purged from acidified samples with N2. Because of these results, the use of sodium azide (NaN3) as a biostat was evaluated. The results from the 24-h studies showed that the distribution of LAS in solution with EPA-12 sediment was unchanged in solutions containing 0.01 M or higher concentrations of NaN3. Except where noted, all of the work reported in this study was conducted in the solutions of 0.01 M NaN3. NaN3 is the salt of a weak acid (pKa ) 4.7) and behaves as a very weak pH buffer under the pH range 6.5-7.5 of the sediment-water solutions used in this study. NaN3 was selected for use since it sorbs very weakly to the sediments. In independent experiments, we demonstrated that sorption of LAS is independent of solution composition, for solutions of composition x M NaCl/(0.01 - x) M NaN3. Some other biostats, such as HgCl2 used by Matthijs and De Henau (19), interact more strongly with surfaces and could influence the distribution of the target compound.
Results Isotherms of C-10, C-12, and C-14 LAS on Sediment EPA12. Sorption isotherms of C-10, C-12, and C-14 LAS on sediment EPA-12 (2.3% organic carbon) are shown in Figure 1. The isotherms are only slightly nonlinear at surface concentrations below 50 µmol/kg but become more nonlinear at higher concentrations. The isotherms can be represented well by either the Freundlich equation (eq 1) or the virial equation (eq 2), for which the parameters are given in Table 3. Effects of pH on the Sorption of C-10, C-12, and C-14 LAS. The effects of [H+] on the sorption of the LAS homologues are shown in Figure 2. Values of log D as a function of log [H+] are shown for the sorption of C-10 LAS, C-12 LAS, and C-14 LAS on sediment EPA-12. The concentrations of LAS were low in all experiments: less than 50 µmol LAS/kg of sorbent and less than 50 nM LAS in solution. Thus, the data were obtained from the near-linear region of the isotherm, and comparison of values of D is straightforward. It is apparent from the data in Figure 2 that the slopes (∆ log D/∆ log[H+]) for all three homologues are approximately equal and that the differences in intercepts are approximately equal. These constraints were imposed through the equation
log Dc ) a log [H+] + bnCH2 + c
(3)
TABLE 3. Sorption Isotherms of LAS Homologues on Sediment EPA-12 in 0.01 M NaN3a Freundlichb
virial/electrostaticc
homologue
log K
n
K, L/kg
b, kg/mol
C, F/m2
C-10d
0.58 0.98 2.27
0.90 0.86 0.93
15.1 77.0 709
2160 2160 2160
0.14 0.14 0.14
C-12e C-14d
a Values of the parameters in the Freundlich equation and the virial equation determined from data in Figure 1. b Parameters determined by unweighted linear regression of the Freundlich equation in log form: log CLAS(s) ) log K + n log CLAS(w), with CLAS(w) in mol/L and CLAS(s) in mol/kg. c Parameters determined by weighted nonlinear regression of the virial equation in the form CLAS(w) exp[-bCLAS(s)]K ) CLAS(s), with the weighting factor based on a constant relative error in CLAS(s). The relation between the parameter b and the capacitance in the electrostatic model is b ) F2/RTCs, as can be derived from the equations in Table 1. d EPA-12 at 0.048 kg/L. e EPA-12 at 0.014 kg/L.
FIGURE 3. Effect of added Ca2+ on the adsorption of LAS. The value of log D (distribution ratio) is plotted as a function of the logarithm of concentration of Ca2+ added to the solution for sorption of C-10, C-12, and C-14 LAS on sediment EPA-12 at 0.0094 kg/L in 0.01 M NaN3. The lines were calculated from eq 4 with the parameters in Table 4.
FIGURE 2. Effect of concentration of H+ on the sorption of LAS. The value of log D (distribution ratio) is plotted as a function of log [H+] values (before and after the equilibration period) for sorption of C-10 LAS, C-12 LAS, and C-14 LAS on sediment EPA-12 at 0.048 kg/L in 0.01 M NaN3. The lines were calculated from eq 3 with the parameters in Table 4.
TABLE 4. Effect of H+ and Ca2+ on Sediment-Water Distribution Ratios of LASa ion
a
b
c
H+ Ca2+
0.173 0.219
0.379 0.342
-1.24 -1.28
a Parameters are determined from data in Figures 2b and 3c for the equation log D ) a log [i] + bnCH2 + c, where [i] ) [H+] or [Ca2+] (mol/L). b EPA-12 in 0.01 M NaN at 0.048 kg/L. c EPA-12 in 0.01 M NaN at 0.0094 3 3 kg/L.
where n is the number of methylene units in the alkyl chain and values of the adjustable parameters a, b, and c were determined by weighted linear least squares; the values of the adjustable parameters are given in Table 4. The dependence of D on the number of methylene groups can be explained in terms of the hydrophobic effect, and the dependence on pH can be attributed either to electrostatic interactions or to specific chemical interactions with surface functional groups, as will be discussed. Effects of Calcium on the Sorption of C-10, C-12, and C-14 LAS. Similar experiments were performed to examine the effects of [Ca2+] on the sorption of LAS homologues; the results are shown in Figure 3. The addition of Ca2+ clearly
FIGURE 4. Sorption isotherms of C-12 LAS on different sediments at 0.024 kg/L in 0.01 M NaN3. The lines were calculated with the virial equation (eq 2) with the parameters in Table 5. promotes sorption. This effect is even slightly greater than that of H+, which is shown in Figure 2. The parameters a, b, and c, which were determined from the fit of the equation
log D ) a log [Ca2+] + bnCH2 + c
(4)
to the data in Figure 3, are given in Table 4. Sorption of C-12 LAS on Different Sediments. A set of experiments was run in an attempt to establish the properties of sediments that influence the sorption of LAS. Isotherms for the sorption of C-12 LAS on four of the five sediments are shown in Figure 4; sorption of LAS to EPA-1 was too low to be determined. The experimental data have been interpreted in terms of a Freundlich equation (eq 1) and the virial equation (eq 2), for which the parameters are given in Table 5. Isotherms of C-12 LAS on EPA-12 at Different Sediment Concentrations. In an effort to clarify the effect of the concentration of solids on the observed concentration distribution ratio, a series of isotherms were obtained with a single surfactant on a single sediment, but with different VOL. 33, NO. 18, 1999 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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TABLE 5. Sorption Isotherms of C-12 LAS on Different Sediments in 0.01 M NaN3 at 0.024 Kg/La Freundlichb
virial/electrostaticc
sediment
log K
n
K, L/kg
b, kg/mol
C, F/m2
EPA-13 EPA-12 EPA-16 EPA-25
1.84 0.85 0.46 0.68
0.92 0.83 0.80 0.85
288 129 90.8 65.1
3 770 1 290 20 000 18 100
0.083 0.24 0.016 0.017
a Values of parameters are determined from data in Figure 4. Parameters determined by unweighted linear regression of the Freundlich equation in log form: log CLAS(s) ) log K + n log CLAS(w), with CLAS(w) in mol/L and CLAS(s) in mol/kg. c Parameters determined by weighted nonlinear regression of virial equation in the form CLAS(w) exp[-bCLAS(s)]K ) CLAS(s), with the weighting factor based on a constant relative error in CLAS(s). The relation between the parameter b and the capacitance in the electrostatic model is b ) F2/RTCs, as can be derived from the equations in Table 1. b
FIGURE 5. Sorption isotherms of C-12 LAS on sediment EPA-12 at different ratios of mass of sorbent to volume of solution (or concentration of sorbent in water, Cs(w)) in 0.01 M NaN3. Lines were calculated from the Freundlich equation (eq 1); the parameters are not given. concentrations of sediments in water. The results are shown in Figure 5 and are partially explained by a model which incorporates the increase in concentration of Ca2+ with the increase in sediment concentration, as will be discussed.
Discussion The primary objective of this study was to understand the sorption of LAS to environmental sorbents in terms of the hydrophobic, specific chemical, and electrostatic interactions. Ideally, we would approach this objective by designing experiments in which one parameter at a time is varied while the others are held constant. However, in practice, this approach is impossible to realize: (i) isotherms are nonlinear, rendering the concentration distribution ratio dependent on the total amount of material in the system; (ii) the properties of the environmental sorbents (e.g., fraction organic carbon, specific surface area, etc.) vary; and (iii) it is impossible to vary one parameter of solution chemistry in the sediment slurry without affecting many others (e.g., a change in pH also changes the free Ca2+ concentration dramatically). Furthermore, the environmental sorbents are themselves heterogeneous. Thus, it is often difficult to distinguish among several possible causes for any particular effect. Energy of Sorption of LAS. The driving force for the sorption of LAS can be analyzed as contributions from three 3114
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primary sources: (i) hydrophobic interactions, (ii) specific chemical interactions of the sulfonate group, and (iii) electrostatic interactions. Even for homogeneous sorbents, such a separation of energies into discrete components is often indeterminate; for heterogeneous natural sorbents, the problem is much more difficult. However, such a mechanistic approach does aid in understanding the sensitivity of sorption to variations in environmental conditions. Hydrophobic Interactions. Hydrophobic interactions literally refer to the tendency of water to expel the solute molecule. Two questions arise with respect to hydrophobic interactions and sorption of LAS: (i) the effect of the alkyl chain length and (ii) the sorption of alkylbenzenesulfonate compared to analogous alkylbenzenes. The effects of alkyl chain length can be seen in Figures 1-3. The difference in log D for sorption of the C-10, C-12, and C-14 homologues is equivalent to ∆ log K ≈ 0.4 per CH2 group. This value is similar to the value of ∆ log K/∆nCH2 ) 0.45, reported by Hand and Williams (20) (as the multiplicative factor 2.8) for the sorption of LAS on a series of environmental sorbents and ∆ log K/∆nCH2 ) 0.44 reported by Dick et al. (21) (as ∆G° of 600 kcal/mol) for the sorption of LAS onto oxides. Furthermore, this value is similar to that of a CH2 group on the solubility of a nonpolar hydrophobic compound in water, as demonstrated by the fragment constant of ∆ log K/∆nCH2 ) 0.50 reported by Lyman (22). The fact that the energy of sorption per CH2 group is similar for many different sorbents is noteworthy. It implies that the energies of interaction of the CH2 group with the different types of sorbents are similar or simply that the differences are small compared to the hydrophobic interaction of CH2 with the common solvent water. The second question involves the contribution of the hydrophobic alkylbenzene and the ionic sulfonate group. The log K for the sorption of dodecylbenzene can be estimated by many methods; in any case, it is several orders of magnitude greater than that of dodecylbenzenesulfonate. The difference in sorptivity between dodecylbenzene and dodecylbenzenesulfonate can be attributed to the tendency of the sulfonate group both to promote the solubility of the dodecylbenzene in water and to repel the molecule from negatively charged surfaces. Interactions of the Sulfonate Functional Group. Four types of results (from this study and the literature) can be interpreted as evidence for interaction of the sulfonate group: (i) increase in sorption of LAS with increase in solution Ca2+ concentration; (ii) increase in sorption of LAS with increase in solution H+ concentration; (iii) correlation of sorption of LAS with sesquioxide content of environmental sorbents (18, 23-25); and (iv) sorption competition between LAS and other anions (SO42- and HPO42-) (18) and sorption competition between sulfate and other anions (e.g., Nodvin et al. (26); Singh (27); Chao et al. (28)). These results could be interpreted as specific chemical interactions of the sulfonate group or as nonspecific electrostatic interactions. It is not easy to distinguish between the two. Dependence of LAS Sorption on Ca2+ Concentration. Two mechanisms could be invoked to explain the increase in sorption of LAS with the increase in Ca2+ concentration (Figure 3). LAS could form a positively charged complex with Ca2+
Ca2+ + RSO3- ) RSO3Ca+
(5)
that would then be preferentially adsorbed on the negatively charged sediment particle; however, the formation constant for the reaction
Ca2+ + SO42- ) CaSO4(aq)
(6)
is small and presumably still smaller for the alkylbenzenesulfonate, for which no formation constants could be found. Alternatively, the Ca2+ could adsorb directly to the sediment particle, reducing the negative surface charge and the electrostatic repulsion of LAS. In separate experiments, we have found that Ca2+ at millimolar concentrations can significantly affect the electrophoretic mobility (and surface charge) of the EPA-12 particles. A third possible mechanism for the effect of Ca2+, the precipitation of Ca(LAS)2 (29), can be ruled out, since the concentrations of LAS in these experiments were on the order of 10-7 M, orders of magnitude below saturation. Dependence of LAS Sorption on pH. The dependence of sorption of LAS on H+ activity is consistent with that of other ionic molecules on the surfaces of environmental particles. There are two basic mechanisms through which this dependency arises: (i) specific pH-dependent reactions, such as surface complexation or ligand exchange, and (ii) the development of pH-dependent surface charge. Inoue et al. (18) have discussed the pH dependence of LAS adsorption in terms of ligand-exchange competition of OH- with LAS through reactions such as
XOH + RSO3- ) XOSO2R + OH-
(7)
where XOH represents a surface hydroxyl group such as tFeOH or tAlOH. They have also shown a correlation of the degree of LAS sorption with Fe- and Al- oxide content of sorbents. The other possible interpretation of the pH dependence is electrostatic interactions caused by pH-dependent surface charge from reactions of surface hydroxyl groups, for example,
XOH ) XO- + H+
(8)
XOH + H+ ) XOH2+
(9)
or surface organic matter, for example,
YCOOH ) YCOO- + H+
(10)
YOH ) YO- + H+
(11)
where Y represents the backbone of the environmental organic matter. Interpretation of the pH Dependence through the Electrostatic Effect. The change in log D of 0.17 unit per unit change in pH (Figure 2) corresponds to a change in surface potential of 9 mV/pH unit, as determined from the factor in Table 1, eq T-5. Coincidentally, the electrophoretic mobility of EPA-12 particles, when converted to the equivalent ζ-potential for spherical particles, also yielded a variation of 9 mV/pH unit (unpublished results). In view of all of the uncertainties associated with ζ-potential determination (30), this correspondence cannot be taken as proof of the electrostatic effect but is taken simply as an indication that it is of the right order of magnitude to explain the results. The origin of the ≈10 mV/pH unit dependence remains unclear. This value is far under that found for many pristine oxide surfaces (45-60 for oxides of Al, Fe, and Mn and 30-45 for Si) but certainly not unexpected, considering the probable low density of ionizable groups and high degree of nonplanarity of the interface. Both of these factors lead to a small change in potential with hydrogen ion activity (31). The deficiencies of a model based on planar interface geometry for explaining the electrostatic energy of sorption to surfaces of environmental particles have been discussed by Westall (31). An alternative to the planar model is the gel or Donnan model. This model has been used by Sakata and Katayama (32, 33) to describe adsorption of anionic and
cationic solutes on synthetic fibers with ionogenic (carboxylic acid) functional groups, not dissimilar to those of natural geopolymers. Whereas this model is probably more realistic than a planar interface model for the sediments in this study, it still does not enable us to represent the gradual effect of pH (∆ log D/∆ log [H+] ) 0.17), over a range of 5 pH units, without invoking surface heterogeneity (i.e., a distribution of surface sites with a spectrum of acidity constants and affinity constants for LAS). Thus, the observed pH dependence of sorption is consistent with the pH dependence of the effective ζ-potential, but one still needs to invoke surface heterogeneity to explain both of these pH dependencies. Nonlinearity of Isotherms. All of the isotherms in this study are nonlinear (i.e., slope < 1 on a log-log scale or simply nonlinear on a linear scale). Nonlinearity can result from many causes, including (i) sorbate-sorbate interactions, such as electrostatic repulsion; (ii) heterogeneity of the sorbent and sequential saturation of sites of decreasing affinity and increasing abundance (34); and (iii) artifacts such as undetected changes in pH, major ion composition, and particle aggregation, which sometime accompany systematic changes in sorbent or sorbate concentration. We have reduced the artifacts by pretreatment of the sorbent and consider the contributions of heterogeneity and electrostatics. First, it is noteworthy that the isotherms of all three homologues (Figure 1) can be fit very well with the virial equation with the same b factor (Table 3). This result is consistent with the affinity of the surface being a homologueindependent function of sorbent loading, as would be expected for an electrostatic interaction or a sulfonate-specific interaction. The values of the virial-equation log K (Table 3) for the homologues vary between 0.35 and 0.45 log units per methylene group, close to the value expected from hydrophobic interaction theory. Next, we consider the interpretation of the exponential term in the virial equation in terms of electrostatic theory. From the equations in Table 1 and the data in Table 3, it can be seen that the b factor would be equivalent to an interfacial capacitance of 0.15 F m-2, if LAS were adsorbed uniformly to a perfectly planar two-dimensional interface (as opposed to three-dimensional absorption), if the sorbent were perfectly uniform in its affinity for LAS, and if the N2 BET surface area were an accurate measure of this two-dimensional surface. These conditions are, of course, somewhat removed from reality, but they are consistent with current usage of electrostatic models in environmental science, and they allow an order of magnitude estimate to be made about the possible contribution of electrostatic effects. For comparison, the absolute minimum for the interfacial capacitance is 0.2 F m-2, as computed from the Gouy Chapman theory for a perfectly planar interface in 0.01 M 1:1 electrolyte at the potential of zero charge (35); theory would predict the capacitance to be still greater for an irregular interface or a charged surface. Thus, it appears that the nonlinearity of these isotherms, even at these low surface concentrations, is not completely inconsistent with an electrostatic origin. Furthermore, the virial equation allows us to assess the degree of nonlinearity of the isotherms. For surface concentrations of LAS of less than 9.4 µmol/kg, the magnitude of the exponential factor is 0.98 e exp(-bCLAS(s)) e 1, or the nonlinearity is less than 2%. Through similar calculations, the error made by assuming a linear isotherm can be approximated. It should be noted that the surface concentrations of LAS in these experiments are far, far below those anticipated for monolayer coverage of LAS. Thus, surface saturation as a cause of nonlinearity, or hemimicelle formation, is not expected in these experiments. A monolayer of cubic packed sulfonate groups (based on the diameter of 0.59 nm for the sulfonate head, given by Somasundaran and Fuerstenau (36)) VOL. 33, NO. 18, 1999 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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on a perfectly uniform surface area of 10 m2 g-1 would result in an sorption density of 4.8 × 10-6 mol m-2 or 48 000 µmol kg-1; the sorption densities in these experiments (Figures 1, 4, and 5) are clearly far below this limit, for uniformly distributed LAS. Finally, it is worth noting that Hand and Williams (20) report linear isotherms for both mixed C-10 LAS and mixed C-14 LAS in experiments run at constant concentrations of sediments in water and roughly the same concentrations of surfactants on the surface and in solution as those used in these experiments. The differences in the linearity of the isotherms in that study and this study are attributed to the differences in the sorbents, as discussed below. Sorption of C-12 LAS on Different Sediments. One experiment was carried out to compare the sorption of LAS on different sediments. Four sediment properties were of particular interest: organic carbon content, extractable iron and aluminum, pH, and surface area, as determined directly or estimated from the size distribution among sand, silt, and clay size classes. The properties of the sediments are listed in Table 2. Organic carbon ranges from 0.2 to 3%, total extractable Fe + Al ranges from 44 to 421 µmol/g, the pH ranges from 6.7 to 7.6, and the size distribution ranges from virtually all sand to virtually no sand. Sorption isotherms of C-12 LAS on these sediments are shown in Figure 4, and the parameters for the Freundlich and virial equations are listed in Table 5. The degree of nonlinearity of the isotherms varies widely. In view of the nonlinearity of these isotherms at higher concentrations, there is no simple, unambiguous way to compare quantitatively the energy of sorption of LAS to the different materials. However, a very simple qualitative comparison can be made. Inspection of the isotherms shows that the distribution of surfactant between the sediment and water, per unit mass of sediment, decreases in the order EPA-13 > EPA-12 > EPA16 ≈ EPA-25 > EPA-1. This order corresponds to the order of organic carbon content of the material and is not inconsistent with the order of extractable Fe + Al. A very weak correlation of sorption energy with organic carbon content of sediment has been seen by Hand and Williams (20) and Urano et al. (37). Several researchers have found a correlation between sesquioxide content (Fe + Al) of the sorbent and adsorption of sulfonate and inorganic ions such as sulfate and phosphate (18, 24). Although the influence of the pH on the distribution of LAS between water and a single sediment is easy to establish, it is clear from these data that the solution pH alone does not dominate the partition behavior. There is virtually no correlation of the order of sorption energies with solution pH and some notable exceptions (EPA-12, for example). Since the fraction of these sediments in the clay size fraction correlates relatively well with the organic carbon content, an assessment of the relative impact of these factors is difficult. A comparison between the isotherms of EPA-16 and EPA25 is particularly interesting. The isotherms are virtually identical, but all factors (organic carbon, Fe + Al, pH, and size fraction) would appear to favor adsorption to EPA-16. A more rigorous comparison of the effects of sediment properties on sorption behavior would require a much broader set of data. Effect of the Ratio of Solids to Liquid in the Slurry. An apparent dependence of the concentration distribution ratio on the ratio of solids to liquid in the slurry has been reported frequently for batch sorption studies (38). Some of the factors that can contribute to this phenomenon are (i) the dependence of solution composition (concentration of NSM or species truly dissolved) on the amount of solids in solution, (ii) the slow sorption or desorption kinetics, (iii) the particle aggregation, (iv) the variation of D with nonlinear isotherms, 3116
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or (v) the use of mixtures of sorbing compounds rather than pure substances. As isotherms are known to be nonlinear in this study, the isotherms themselves, at various ratios of solid to liquid in the slurry, are compared in Figure 5. Although many explanations could be invoked to explain the differences among the isotherms, we begin with one that is to a great extent independently verifiable: the concentration of Ca2+ in solution varies with the amount of solid, and the value of D varies with the concentration of Ca2+. We have shown (Figure 3) that ∆ log D/∆ log [Ca2+] ≈ 0.23 for C-12 LAS and EPA-12 at concentrations of LAS on the sediment of less than 30 µmol/kg. In other experiments, we determined the concentration of Ca2+ as a function of the amount of EPA-12 in 0.01 M NaN3: we found ∆ log [Ca2+]/∆ log Cs(w) ≈ 0.66, for EPA-12 in 0.01 M NaN3 over the range of Cs(w) from 0.005 to 0.01 kg/L. This variation was more or less consistent with ion exchange between Na+ from solution and Ca2+ from the surface, as determined from charge balance in solution. From these data, we can calculate the variation in D with Cs(w) that is attributable to the variation in [Ca2+]: 2+ ∆ log D ∆ log D ∆ log [Ca ] ) ) 0.18 (12) ∆ log Cs(w) ∆ log [Ca2+] ∆ log Cs(w)
Finally, we estimated the variation in apparent D among the different isotherms in Figure 5: we defined Deff as the slope of the Freundlich isotherm evaluated at C(s) ) 20 µmol/kg
Deff ) dC(s)/dC(w)
(at 20 µmol/kg)
(13)
s(w) ) 0.18, which coincides with the value determined from [Ca2+] dependence. This good agreement is perhaps somewhat fortuitous, but at least it shows that the [Ca2+] dependence is of the right order of magnitude to explain the observed particle concentration dependence of log D. (Note that, under these experimental conditions, the value of Deff increases with Cs(w), in contrast to the “classical solids effect” in which D decreases with Cs(w).) The definition of Deff in eq 13 is based on the lowconcentration, near-linear segments of the isotherms. At higher concentrations, the variation of Deff with Cs(w) is stronger. To investigate these regions, more experiments would have to be performed with very close control of solution composition. Furthermore, a conceptual model for the nonlinearity of the isotherms would have to be invoked. The point to be made is that many mechanisms can contribute to the apparent dependence of D on Cs(w). To interpret this effect mechanistically requires the utmost attention to all experimental conditions. Comparison of Calculated and Predicted Values of log K. Di Toro et al. (39) have proposed a method to estimate the sorption of anionic surfactants on sediments based on four principal factors: the critical micelle concentration (cmc) as a property of the surfactant, the fraction organic carbon (foc) or the cation-exchange capacity (CEC) as properties of the sediment, and the concentration of solids, as a property of experimental conditions to account for manifestations of the classical solids effect (i.e., decreases in D with Cs(w)). The empirical model of Di Toro et al. was calibrated for the lowconcentration, linear range of isotherms reported in the literature; thus, their predicted distribution ratios should correspond to the virial K (Tables 3 and 5) discussed in this paper. The observed and predicted values of log K for C-12 LAS on four sediments are presented in Table 6. The observed value is simply the virial K transcribed from Table 5. Two predicted values are presented: one that represents the “low solids concentration, no solids effect” limit in experimental
and found that ∆ log Deff/∆ log C
TABLE 6. Comparison of Observed Values of the Virial log K for C-12 LAS on Four Sediments to Those Predicted by the Model of Di Toro Et al. (39) predicted sediment
obsd log K,a L/kg
log K,b L/kg
log Kc, L/kg
EPA-13 EPA-12 EPA-16 EPA-25
2.46 2.11 1.96 1.81
3.33 3.27 3.15 3.06
1.62 1.62 1.61 1.61
a From virial equation, Table 5. b log K (no “solids effect”) calculated from the model of Di Toro et al. (39); the value of K in this column corresponds to their πc calculated from their eq 8 (πc ) c1focc2/[cmc], where C1 and C2 are constants, foc is the fraction of organic carbon, and [cmc] is the critical micelle concentration) with parameters in their Table 8 and [cmc] calculated from their eq 3 (ln [cmc] ) -8.038 mol/L); values of foc are from Table 2 of this paper. c log K (“corrected” for “solids effect”) according to model of Di Toro et al. (39): the value of K in this column corresponds to their π calculated from their eq 7 (π ) πc/(1 + mπcνx), where m is the ratio of solids to liquid, πc is defined in footnote b, and νx is a parameter) and eq 8 (defined in footnote b) with parameters in their Table 8, and cmc calculated from their eq 3 (ln cmc ) -8.038 mol/L); values of foc are from Table 2 of this paper, and the ratio of solids to liquid is 0.024 kg/L.
design (i.e., based only on cmc and foc) and one that includes explicitly an adjustment for the “solids effect” (i.e., based on cmc, foc, and solid-to-liquid ratio). The one that is adjusted for the solid-to-liquid ratio is almost independent of foc, with partitioning predicted to be dominated by the solids effect. The predicted values of log K span the observed values of log K, but agreement is not very good. As can be seen from a comparison of the columns in Table 6, a particular problem is how to handle the solids effect. As stated above, many factors can logically contribute to this effect, including the presence of nonseparable suspended material, the concentrations of which increase with the solid-to-liquid ratio (16); the presence of radiochemical impurities or degradation products in experiments conducted with radiolabeled compounds (40); and the use of mixtures of compounds (e.g., surfactant mixtures) instead of the pure substances. Then the question is, is it better to clean up the experimental design and remove the solids effect to obtain “intrinsic” distribution ratios or is it better to have “operational” distribution ratios, which may or may not reflect the conditions of the environment that one is trying to analyze. The answers to these questions are beyond the scope of this paper, but they illustrate one of the difficulties encountered as one strives for universally applicable correlations. Environmental Behavior of LAS. The long-range goal of this mechanistic study of sorption is to gain understanding that can be used to elucidate the environmental behavior of LAS. We have shown that the sediment-water distribution of LAS depends on solution chemistry ([H+] and [Ca2+]), sorbent composition (foc and extractable Fe + Al), and LAS homologue (hydrophobic interaction) and that some aspects of the observed particle concentration effect can be explained by co-variations in solution chemistry. For design and interpretation of controlled laboratory experiments, this level of detail can be significant. However, for application to field systems, some simplifications might be justifiable. For example, one might simply be aware of the sensitivity of LAS distribution to solution chemistry but not model it explicitly, using operational constants instead. Toward the goal of a priori prediction of LAS distribution based on sorbent properties and LAS structure, one might reach the conclusion from the data presented here that an equation as reliable as that for hydrophobic partitioning for neutral molecules will never be developed for these amphiphilic molecules such as LAS. However, if the distribution is measured for one homologue, distributions of others could
be predicted from hydrophobic theory or the relative critical micelle concentration (39). Of course, the level of approximations that is justifiable always depends on the purpose for which the prediction is to be used.
Acknowledgments We gratefully acknowledge Keith Booman, Richard Sedlak, and Alvaro DeCarvalho of the Soap and Detergent Association and Vincent Hand now of Miami University for their contributions to this work. This research was sponsored by the Soap and Detergent Association.
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Received for review April 27, 1998. Revised manuscript received January 6, 1999. Accepted January 20, 1999. ES9804316
