Environ. Scl. Technol. 1991, 25, 1585-1589
Removal of Sparingly Soluble Organic Chemicals from Aqueous Solutions with Surfactant-Coated Ferrihydrite Thomas M. Holsen,” Elalne Ruth Taylor, Yong-Chan Seo, and Paul R. Anderson
Pritzker Department of Environmental Engineering, Illinois Institute of Technology, Chicago, Illinois 606 16
A surfactant, sodium dodecyl sulfate (SDS), was adsorbed onto ferrihydrite at low pH and desorbed at high pH. Adsorption isotherms of SDS on ferrihydrite yielded an S-shaped curve typical for surfactant adsorption on mineral surfaces. This behavior is due to the surfactant forming aggregate structures on the solid surface. SDS adsorption and desorption experiments revealed that both processes were fast, with equilibrium being reached within 1 h. Toluene, p-xylene, and trichloroethylene sorption experiments demonstrated that the SDS-coated ferrihydrite was able to remove these sparingly soluble organic chemicals (SSOCs) from solution. The amount of SSOC sorbed was directly related to the amount of SDS coating the ferrihydrite. The SSOC with the lowest Solubility (p-xylene) had higher removals than the high-solubility SSOC (TCE). Introduction There has been a growing interest in the fate of the enormous quantity of surfactants that are used in industrial and household applications because they could influence the behavior of other pollutants in surface water or groundwater (I, 2). Recent studies have shown that both surfactants and dissolved organic matter (DOM) enhance the aqueous solubility of sparingly soluble organic compounds (SSOCs) (1-6). This water-solubility enhancement is a result of a partitioning of the solute into the organic environment of the high molecular weight DOM (3,4)or, in the case of surfactants above their critical micelle concentration (CMC),into the organic environment provided in the interior of the micelle (I, 2, 5, 6). The amount of solubility enhancement has been found to increase with decreasing solute solubility. In addition to interactions with each other, both DOM and ionic surfactants can adsorb to charged surfaces in the environment. Early workers in the field of surfactant adsorption found that, at low concentrations of surfactants and at low surface potentials, surfactant ions are adsorbed, primarily due to electrostatic forces between the charged surface and the oppositely charged surfactant ion (7,B). This one-to-one charge exchange of surfactant and adsorbed counterions does not significantly change the electrophoretic mobility (charge) of the solid particle (zone 1,Figure 1). At higher surfactant solution concentrations, aggregates begin to form as the hydrophobic tails of the adsorbed surfactants begin to associate. This association greatly enhances the adsorption of surfactant, resulting in a rapid rise in the adsorption isotherm (zone 2). Under these conditions adsorption is due to both electrostatic attraction of the surface for the oppositely charged surfactant ion and hydrophobic attraction between the surfactant tails. Furthermore, the electrophoretic mobility of the particle rapidly decreases as the surface charges are neutralized. When the surface of the particle is completely neutralized, the electrostatic attraction disappears, and adsorption levels off because further association is due only to hydrophobic interactions with the surfactant tails (zone 3) (7, 8). 0013-936X/91/0925-1585$02.50/0
In a recent paper, Yeskie and Harwell (9) presented theoretical calculations regarding the adsorption of surfactants on charged surfaces and slightly revised the earlier model. The purpose of their study was to examine the relative importance of hydrophobic interactions compared to electrostatic interactions in determining the surfactant structure on a surface. The model considered the formation of a monolayer of surfactant (a hemimicelle) composed of surfactant molecules adsorbed to the surface with their hydrophobic tails sticking out into solution and of a bilayer of surfactant (an admicelle), which forms when additional surfactant molecules sorb “tail first” to the hemimicelle extending their hydrophillic heads into solution. Results from their modeling suggest that electrostatic interactions govern the transition from the hemimicelle structure to the admicelle structure. The admicelle structure becomes favored at higher charge densities away from the zero point of charge of the mineral surface. The authors pointed out that if the charge distribution on the mineral surface on which the surfactant is adsorbing is heterogeneous, systems could exist in which portions of the surface are covered with admicelles while other portions of the surface are covered with hemimicelles. Their modeling results also supported earlier experimental work by Bisio et al. (IO) on surfactants adsorbed to a polycarbonate surface. They found that the tail groups of the outer layer of the bilayered aggregate do not significantly interpenetrate the tail group region of the inner layer. The realization that adsorbed surfactants will provide an organic coating on charged particles allowing SSOCs to partition into them has lead some researchers to coat low organic carbon containing clays with surfactants, greatly increasing their ability to remove SSOCs from solution (11). This coated clay could then be used in landfill liners where both low permeability and high SSOC sorption capacity are desired. The research presented in this paper explores the adsorption of an anionic surfactant, sodium dodecyl sulfate (SDS), onto ferrihydrite and the ability of this modified ferrihydrite to sorb SSOCs from solution. The study shows that SDS can be adsorbed by ferrihydrite at low solution pH, enabling the composite adsorbent to remove toluene, trichloroethylene, and p-xylene from solution. When the solution pH is increased, both the surfactant and SSOC are desorbed. The amount of SSOC removed from solution was found to be dependent on its aqueous solubility and the amount of SDS adsorbed onto the ferrihydrite. Experimental Section Sodium dodecyl sulfate (SDS) was purchased from Sigma Chemical Co. Ferric nitrate nonahydrate, sodium nitrate, chloroform Optima, pesticide grade toluene, and pesticide grade p-xylene were purchased from Fisher Scientific. Pesticide grade trichloroethylenewas purchased from Aldrich. All chemicals were used as received. NANO-pure water was used in all of the experiments, and 1.0 M NaOH and HN03were used for pH adjustment. After equilibration of surfactant and ferrihydrite, a final
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Environ. Sci. Technol., Vol. 25, No. 9, 1991 1585
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solution pH was determined, the sample was centrifuged, and the supernatant was removed for analysis. All experiments were performed in a 0.12 M NaN03 swamping electrolyte. Ferrihydrite was prepared by raising the pH of a stirred solution of ferric nitrate, sodium nitrate, and NANO-pure water from 2.0 to approximately 7.5. The solution was allowed to age for a minimum of 2 h before use. Typical concentrations of ferric nitrate and sodium nitrate were 10-2.6and M, respectively. One of two methods was used for SDS analysis, depending on the surfactant concentration. At low concentrations standard methods 512 B (12),anionic surfactants as methylene blue active substances (MBAS), was used. A t high concentration (greater and equal to M) the surfactant concentration was measured with a TOG analyzer (Dohrmann Model DC-80). This method gave results comparable to standard methods 512 B at these concentrations but was much faster. Adsorption isotherms and the affect of pH on adsorption (adsorption edge) were measured with batch experiments as follows. A suspension of ferrihydrite was adjusted to pH values between 4.5 and 5.0 for isotherm experiments and between 8 and 8.5 for adsorption edge experiments. A known amount of SDS was added, the suspension was stirred for 10 min, and a 20-mL sample was collected. For the isotherm experimenh this entire process was repeated to until the SDS concentration increased from M. For the adsorption edge experiments the SDS concentration was held constant and the pH was incrementally decreased. In all cases, the samples were placed on a tumbler to equilibrate for at least 2 h. The amount of SDS adsorbed was determined from the difference between the total amount of SDS added to the suspension and the amount remaining in solution. Control experiments showed that the amount of SDS adsorbed by the glassware was negligible. In adsorption kinetics experiments the SDS-ferrihydrite suspension pH was not adjusted but remained relatively constant between 5.0 and 5.4. Samples were withdrawn at specified times and centrifuged immediately, and the supernatant removed and analyzed. SSOC Sorption. Organic chemical concentrations were measured with a gas chromatograph (GC; Varian Model 3700) equipped with a Supelco SP-1500 column and FID detector and Hewlett-Packard Model 19395A headspace analyzer. The procedure described earlier for ferrihydrite preparation was followed for these experiments except that 1588 Environ. Scl. Technol., Vol. 25, No. 9, 1991
the initial ferrihydrite concentration was higher. After the SDS was added to the ferrihydrite solution, the suspension was aged for a minimum of 2 h (but usually overnight) at pH 4.5. Concurrently, two to five pH-adjusted solutions, each containing 10-3.4M SDS and a varying amount of SSOC were prepared and stirred a minimum of 1 2 h to ensure SSOC dissolution. The following day, the surfactant-coated ferrihydrite was combined with an SSOC-SDS solution in duplicate headspace vials. Duplicate samples were then put in a tumbler for a minimum of 4 h before the samples were analyzed by headspace analysis. The difference between the headspace SSOC concentration in vials with and without ferrihydrite (at equal equilibrium solution surfactant concentration) was used to calculate the mass of SSOC sorbed. Blank experiments with ferrihydrite and toluene showed that there was no sorption of toluene on uncoated ferrihydrite. The kinetics of toluene sorption was determined by M SDS solution to a ferrihydrite suspension adding a M [FeIT). This mixture (pH 4.5) was allowed to equilibrate before it was diluted with a toluene-SDS solution containing 170 mg/L toluene. Samples were equilibrated on the tumbler, removed, and analyzed over a range of 30 min to 24 h. Results and Discussion SDS Adsorption on Ferrihydrite. The amount of SDS adsorbed by ferrihydrite increased nonlinearly with increasing SDS concentration (Figure 2). At a pH below 7 the isotherms are typical of those for anionic adsorption to a charged surface. An increasing amount of SDS was adsorbed at lower pHs. At pH 7, the isotherm begins with a shallow slope (see inset, Figure 2) but becomes markedly steeper at solution concentrations above 0.005 M. This type of behavior is indicative of hemimicelle or admicelle formation (7,8)and is shown as a transition from zone 1 to zone 2 in Figure 1. Above approximately 0.004 M or half of the CMC M (13)), the isotherm levels off as the hydrophobic forces of aggregate formation compete with attractive forces on the surface. The lack of a clear transition between zones one and two at low pH (
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indicated that the amount of toluene sorbed decreased by approximately 30%. Further evaluation of the data, however, reveals that this is an indirect result of the effect of pH on SDS sorption. When the amount of toluene sorbed was normalized by the amount of SDS adsorbed, the pH of the solution actually had little effect on toluene removal (Figure 7). The sorption of trichloroethylene and p-xylene by the modified ferrihydrite also yielded linear isotherms. The amount of SSOC sorption followed the order p-xylene > toluene > trichloroethylene in order of decreasing hydroEnviron. Sci. Technol., Vol. 25, No. 9, 1991
1587
Table I. Calculated Partition Coefficients between D-Xylene. Toluene. and TCE on Surfactant-Coated Ferrihydrite SDS concn, M 10-2.6
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modified ferrihydrite medium than into a soil of comparable organic carbon content. It is likely that this difference is due to the characteristics of the adsorbents. Soil organic matter is a mixture of both polar and nonpolar functional groups, rendering it a less efficient partitioning medium then the hydrophobic SDS coating.
Measured percent organic c a r b o n
Figure 8. Relationshlp between the measured partltlon coefficient and the measured organlc carbon content of p-xylene, toluene, and TCE on SDS-coated ferrlhydrite.
phobicity as measured by both solubility and Kow. The slopes of the isotherms for all three SSOCs are shown in Table I. There has been a great deal of research on the sorption of SSOCs on soils which has shown that this process is primarily due to the partitioning of the SSOC into the soil organic matter (15). The amount of organic chemical sorbed by a soil can be predicted from the soil’s organic carbon content and the SSOC’s solubility or octanol-water partition coefficient. This same approach was used to predict the sorption of SSOCs by the modified ferrihydrite. The partition coefficient (Kd)for the SSOCs is a linear function of the organic carbon content of the modified ferrihydrite (Figure 8). For each SSOC, the slope of this line is K,, the SSOC-organic carbon partition coefficient (Table 11). Because empirical relationships have been established between K, and the solubility of SSOC (log K , = 3.95 0.62 log S (S, aqueous solubility, mg/L) (16)),the theoretical organic carbon content of the sorbent can be predicted and compared to its measured value (Table I1 and Figure 9). Interestingly, the theoretical organic carbon content of modified ferrihydrite is 2-3 times higher than the calculated carbon content based on SDS adsorption, indicating better partitioning of the SSOCs into the 1588 Envlron. Scl. Technol., Vol. 25, No, 9, 1991
mg/L
Figure 9. Comparison between measured K, values and those predicted from Hassett et al. (76) for soil-SSOC systems.
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Conclusions The adsorption of SDS on ferrihydrite yields an Sshaped curve typical of the adsorption of surfactants on charged surfaces. Adsorption increases as the SDS molecules form aggregates on the ferrihydrite surface. As the SDS concentration exceeds the CMC, adsorption levels off as micelle formation begins to occur in the bulk solution. Consistent with anionic adsorption at oxide surfaces, the highest amount of adsorption of the anionic surfactant is at low solution pHs. As pH increases, the percent adsorption decreases to negligible amounts at ferrihydrite’s zero point of charge. The hydrophobic surface of SDS-coated ferrihydrite is an excellent environment for sorbing SSOCs. Sorption occurs within the first 30 min, and there is a linear relationship between the amount of SSOC sorbed and its aqueous concentration. This linear relationship is similar to SSOC sorption by soils and is typical of hydrophobic sorption of SSOCs from solutions into nonpolar sorbents. This research demonstrates that ionic surfactants at low concentrations can adsorb onto mineral surfaces, rendering them hydrophobic. This coating is a highly efficient partitioning medium for SSOCs and under some environmental conditions can significantly retard SSOC mobility in the environment. Registry No. SDS, 151-21-3;TCE, 79-01-6; toluene, 10888-3; p-xylene, 106-42-3; ferrihydrite, 39473-89-7.
Literature Cited (1) Chiou, C. T.; Kile, D. E. Enuiron. Sci. Technol. 1989,23, 832. (2) Kile, D. E.; Chiou, C. T.; Helburn, R. S. Enuiron. Sci. Technol. 1990, 24, 205. (3) Chiou, C. T.; Malcolm, R. L.; Brinton, T. L.; Kile, D. E. Environ. Sci. Technol. 1986, 20, 502. (4) Chiou, C. T.; Kile, D. E.; Brinton, T. I.; Malcolm, R. L.; Leenheer, J. A.; MacCarthy, P. Enuiron. Sci. Technol. 1987, 21, 1231. (5) Valsaraj, K. T.; Gupta, A.; Thibodeaux, L. J.; Harrison, D. P. Water Res. 1988, 22, 1173. (6) Kile, D. E.; Chiou, C. T. Enuiron.Sci. Technol. 1989,23, 833. (7) Somasundaran, P.; Fuentenau, D. W. J.Phys. Chem. 1966, 70, 90.
Environ. Sci. Technol. 1991, 25, 1589-1596
(8) Somasundaran,P.; Middleton, R.; Viswanathan, K. V. In StructurallPerformance Relationships in Surfactants; Rosen, M. J., Ed.; ACS Symposium Series 253; American
Chemical Society: Washington, DC, 1984; Chapter 17.
(9) Yeskie, M. A,; Hawell, J. H. J. Phys. Chem. 1988,92,2346. (10) Bisio, P. D.; Cartledge, J. G.;Keesom, W. H.; Radke, C. J. J . Colloid Interface Sci. 1980, 78, 225. (11) Lee, J. F.; Crum, J. R.; Boyd, S. A. Environ. Sci. Technol. 1989, 23, 1365. (12) Standard Methods for the Examination of Water and Wastewater, 16th ed; Greenberg, A. E., Trussel, R. R.,
Clesceri, L. S., Ed.; American Public Health Association (APHA): Washington, DC, 1985. (13) Vold, R. D.; Vold, M. J. Colloid and Interface Chemistry; Addison-WesleyPublishing Co. Inc.: Reading, MA, 1983.
(14) Anderson, P. R.; Benjamin, M. M. Environ. Sci. Technol.
1990, 24, 692.
(15) Chiou, C. T. Reactions and Movement of Organic Chemicals in Soils. SSSA Spec. h b l . 1989, No. 22. (16) Hassett, J. J.; Banwart, W. L.; Griffen, R. A. In Environment and Solid Wastes; Francis, C. W., Auerback, S. I., Eds.; Butterworth Woburn, MA, 1983; Chapter 15.
Received for review November 27, 1990. Revised manuscript received April 22,1991. Accepted May 14,1991. This work was supported in part through the Industrial Waste Elimination Research Center, a national research center sponsored in part by the U.S. Environmental Protection Agency Office of Exploratory Research.
The Primary Reaction in the Decomposition of Ozone in Acidic Aqueous Solutions Knud Sehested," Hanne Cotfltzen, Jerzy Holcman, Chrlstlan H. Flscher,t and Edwln J. Hart*
Environmental Science and Technology Department, Ris0 National Laboratory, DK 4000 Roskilde, Denmark The primary step in the thermal decomposition of acidic aqueous ozone solutions is the dissociation reaction: O3 F? 0 + 02.This reaction is supported by the retarding effect of O2 on the decomposition rate, by our activation energy measurements, E A = 79.5 f 8.0 k J mol-', and by isotopic exchange experiments with 18J802. The rates of ozone decomposition were measured at pHs 0-4, at temperatures 0-46 "C, and at O2concentrations 10-5-10-3 M. The 0 atom formed in the dissociation reaction is postulated to be the precursor for OH and H 0 2formation. The ionization of the H 0 2 free radical (pK = 4.8) accounts for the increase in rate with rising pH. The mechanistic implications of O3 decomposition in acid solution are briefly discussed. Introduction The mechanism of aqueous ozone decomposition has been studied for over half a century (e.g., refs 1-23). Early work has shown that ozone decomposes faster in increasingly alkaline solution. The chain-reaction characteristics of the self-decomposition reaction of aqueous ozone have clearly been pointed out, and it is generally accepted that O3 decomposition is base-catalyzed. Previous work (10, 12,23,24)has established that in the presence of radical scavengers the initiation rate above pH 8-9 is proportional to the concentration of both O3 and OH-. One of the proposed initiation reactions in alkaline solutions is
O3 + OH-
-
H02- + O2
(1)
The rate constant of initiation reaction 1 is 40-70 dm3 mo1-ls-l (2,9,10,12), and it has been shown that OH, 02-, and 0, radicals are intermediates in the subsequent chain decomposition. Pulse radiolytic investigations (13-20)have recently elucidated new facets of the propagation and termination steps of the aqueous ozone decomposition by determining the rate constants of the free-radical reactions with ozone. In acid solutions, however, reaction 1cannot be the sole initiation reaction, since the observed reaction rates are +
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0013-936X/91/0~25-1589$02.50/0
much higher than would be predicted from the low OHconcentrations below pH 4. Although an additional initiation reaction is necessary, definite identification of such a reaction has not been forthcoming. Our objective is to identify the primary decomposition reaction in acidic ozone solutions. An early proposal (23) postulated formation of OH radicals in a reaction with water: O3 + H 2 0 20H + O2 (2) This reaction has also been suggested in a more recent study (6). Experiments (25) with 180-enriched water in 0.04 M HCIOl have established little or no exchange between the oxygen atoms of O3 and water, but upon decomposition of O3 an exchange does take place. From the extent of exchange it was concluded that the active intermediate was the OH radical and not H02 or C104. In the present paper we report on the effect of O3 and O2 concentrations and temperature on the thermal decomposition of acidic ozone solutions in the pH range 0-4 and in the temperature range 0-46 "C. Tracer experiments with added 18Ja02 were also carried out in order to determine the importance of the 0 atom as a major initiating species.
-
Experimental Section Special care was taken to avoid the effects of impurities on the decomposition of ozone solutions. All glassware was baked at 450 "C for 4 h after cleaning and then flushed with the solution before use. The water used was triply distilled, and in most experiments preozonized. Other chemicals were reagent grade and were used without further purification. All experiments were carried out with reagent grade HC104. The pH was measured on a Radiometer PHM22s pH meter. The preparation of (2-4) X M O3 stock solutions, nearly free of 02,has been previously described (15-17). Stock solutions were also prepared from O3 gas stored in the dark in moist acidic 100-mL glass syringes at 0 "C. We found that O3 stored as gas under these conditions was much more stable than 0, dissolved in acidified aqueous solutions. However, in order to prevent explosionbf the O3 in the syringes, unusual handling precautions are required. We have found that an explosion of 50-75 mL of
@ 1991 American Chemical Society
Environ. Sci. Technol., Vol. 25, No. 9, 1991
1589
