Environ. Sci. Technol. 1997, 31, 1674-1679
Thermodynamics of Extraction of Copper(II) from Aqueous Solutions by Chelation in Supercritical Carbon Dioxide JENNIFER M. MURPHY Department of Civil and Environmental Engineering, Environmental Engineering Program, University of Connecticut, Storrs, Connecticut 06269 CAN ERKEY* Department of Chemical Engineering, Environmental Engineering Program, University of Connecticut, Storrs, Connecticut 06269
The primary goal of this research was to investigate the feasibility of removing heavy metals from wastewater by chelation in supercritical carbon dioxide. Copper was used as the model contaminant. Extraction efficiency was determined as a function of temperature, pressure, initial moles of chelating agent, and initial moles of copper in a 300-cm3 autoclave. Results indicated that 60% removal of copper can be achieved in a single equilibrium stage with an initial copper concentration of 100 ppm and hexafluoroacetylacetone (HFA), present in excess, as the chelating agent. Experimental extraction efficiencies ranged from 14 to 60% and increased with increasing initial chelate concentration and CO2 density and with decreasing initial copper concentration. A thermodynamic model based on combined reaction and phase equilibria was developed for the prediction of extraction efficiencies. Five aqueous-phase reactions and four phase equilibrium relations were considered in the development of the model. Equilibrium constants and phase distribution coefficients were either drawn from the available literature or measured. The distribution coefficient of the copper-chelate complex was estimated as the ratio of the solubility of the complex in the supercritical phase to its solubility in the aqueous phase. The model successfully predicted the effects of temperature, pressure, and initial amounts of chelating agent and metal contaminant on extraction efficiencies without the use of any adjustable parameters.
Introduction Elevated concentrations of heavy metals in industrial effluents have in recent years become a target for strict legislation. Untreated wastestreams threaten the integrity of water resources, prompting the establishment of limits on metal levels in wastestreams released to the environment. As a result, the demand for cost-effective and environmentally benign metal removal technologies has risen. Several removal technologies have been developed or are in the development stages to treat contaminated effluents. Some examples are hydroxide and sulfide precipitation, ion exchange, membrane processes, and electrochemical opera* Author to whom correspondence should be addressed. E-mail:
[email protected]; phone: (860)486-4601; fax: (860)486-2959.
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tions. Solvent extraction is another technology that may be used to remove metals from wastewater (1, 2). In this technique, the wastewater is mixed with an organic phase consisting of a chelating agent dissolved in a suitable solvent. The pH of the aqueous phase is adjusted to chelate a specific metal. The metal ions react with the chelating agent to form a metal-chelate complex that partitions between the organic and aqueous phases. While it has been established that metal ions can be effectively removed from aqueous systems by solvent extraction, this technique has a few disadvantages. In particular, the ratio of the aqueous-phase volume to the organic-phase volume cannot usually be higher than about 10, which leads to the use of large volumes of solvent, especially when the feed is lean. The fact that significant quantities of solvent are lost by entrainment reinforces this limit. The large consumption of potentially harmful organic solvents in solvent extraction can be eliminated by the substitution of nontoxic supercritical fluids (SCFs). Furthermore, in solvent extraction of metals, the chemical reactions occurring at the interfacial plane may be fast as compared to mass transfer processes. Thus, depending on hydrodynamic conditions in the extraction vessel, the observed kinetics of removal may be controlled by mass transfer. Since the mass transfer characteristics of SCFs are excellent as compared to those of liquid solvents because of their relatively low viscosities and high solute diffusivities, the use of SCFs in place of organic solvents may enhance extraction kinetics. In addition, SCFs generally have a much lower surface tension than traditional organic solvents, resulting in an increased dispersed phase surface area that may reduce the size of equipment required for a particular solvent to feed ratio. Among the SCFs, the nontoxic nature and relatively low critical temperature (31 °C) and pressure (73.8 bar) of carbon dioxide make it an attractive choice for this application. Other notable advantages of supercritical carbon dioxide (SCCO2) are that it is non-flammable, relatively inexpensive, and readily available commercially. Also, the fact that the solvency characteristics of SCFs can be varied with small changes in temperature and pressure may be exploited in the development of selective extraction schemes. As a result of these favorable solvency properties of SCFs, some research and development work has been conducted in various laboratories on the removal of heavy metals from wastewater streams by chelation in SCCO2. These studies began with the pioneering work of Laintz et al. (3), who investigated the extraction of copper(II) from an aqueous solution by chelation with bis(trifluoroethyl)dithiocarbamate using a dynamic extraction scheme. Nearly 100% of the metal was removed from the aqueous sample. The authors indicated that the use of the fluorinated chelating agent yielded much better extraction results than the use of a nonfluorinated analogue, diethyl dithiocarbamate. This effect was attributed to the higher solubility of the fluorinated metal-chelate complex in SCCO2 (4). A trend of increasing extraction efficiency with temperature and density was also noted. Wang and Marshall performed extensive solubility measurements of metal-chelate complexes in SCCO2 (5). They also investigated the removal of zinc, cadmium, and lead from aqueous solutions using tetrabutylammonium dibutyl dithiocarbamate as the chelating agent and SCCO2 as the mobile phase. Nearly complete metal extractions were achieved. Recently, uranyl and Th(IV) ions were extracted from synthetic aqueous solutions and mine waters using thienoyltrifluoroacetylacetone as the chelating agent with extraction efficiencies ranging from 38 to 90% (6). The extraction
S0013-936X(96)00519-6 CCC: $14.00
1997 American Chemical Society
FIGURE 1. Model supercritical chelation system. efficiencies were significantly enhanced by the addition of tributyl phosphate to the system. The same chelating agent was also used to extract trivalent lanthanides from acidic solutions with extraction efficiencies ranging from 17 to 92% depending on the type of metal and the percentage of tributyl phosphate in the SCCO2 phase (7). The results indicated that SCCO2 has the potential to replace commonly used organic solvents in nuclear fuel processing. It is evident from Figure 1, which shows the equilibria involved in the transfer of copper from an aqueous phase to a SCCO2 phase, that chelation is a complex process. The chelating agent, a weak acid, is generally referred to as HA. The distribution of the metal between the aqueous and SCCO2 phases is affected by the equilibrium constants of the aqueousphase reactions and the distribution coefficients of the molecular species. Both the reaction equilibrium constants and the distribution coefficients are theoretically dependent on temperature, pressure, and composition. Chelation in SCCO2 is more complicated than in organic solvents due to the formation of carbonic acid and its derivatives. As seen in Figure 1, the solubility of the metal-chelate complex in the SCCO2 phase represents an upper thermodynamic limit for extraction. However, a high complex solubility in the supercritical phase does not always translate into a high extraction efficiency. The system composition, temperature, pressure, and pH must also allow for sufficient ionization of the chelating agent and formation of the metal-chelate complex. In general, the thermodynamics of the entire system must be favorable for good results. Successful implementation of this technology on an industrial scale for a particular wastestream depends, to a
great extent, on the ability to model and predict the thermodynamic behavior of extraction systems. Since there are currently no such models, this study provides a detailed investigation to determine whether the extraction efficiencies can be interpreted in terms of the phenomena given in Figure 1. Copper was chosen as the model contaminant, and acetylacetone (AA) and hexafluoroacetylacetone (HFA) were tested as chelating agents. System selection was based on the availability of the input data necessary to model the system behavior. Solvent extraction of copper from aqueous solutions using carbon tetrachloride as a solvent was investigated in detail by Sekine and Ihara (8). The authors reported values of copper complex stability constants for acetylacetone and hexafluoroacetylacetone. The mole fraction solubility of cupric acetylacetonate in SCCO2 was reported by Cross et al. (9) to range from 5.084 × 10-6 to 2.582 × 10-5 for temperatures between 35 and 55 °C and carbon dioxide densities between 0.65 and 0.90 g/mL. For cupric hexafluoroacetylacetonate, the mole fraction solubility in SCCO2 was reported by Lagalante et al. (10) to range from 225.1 × 10-5 to 569.9 × 10-5 at 40 °C for pressures between 102 and 340 atm. With knowledge of the solubility limit of the metal-chelate complexes in SCCO2, experiments were conducted at copper concentrations well below this limit. The maximum copper concentration used in the experiments was 1000 ppm in the water phase, and the experiment was performed at 45 °C and 136 atm. The mole fraction solubility of Cu(HFA)2 in SCCO2 at these conditions is approximately 273 × 10-5. If all of the copper in the system were chelated and extracted into the SCCO2 phase, the mole fraction of copper complex in the supercritical phase would be 49 × 10-5, which is well under the solubility limit. With much of the necessary model input data available, the effects of temperature, pressure, initial metal concentration, and initial chelating agent concentration on extraction efficiencies were investigated. The trends indicated by the experimental data are very closely corroborated by a thermodynamic model of the extraction system based on Figure 1. The material presented here should lay the groundwork for further laboratory-scale studies in this area and accelerate development of industrial-scale metal ion extraction systems.
Experimental Section Extraction runs were performed using the apparatus shown in Figure 2. The batch extraction vessel was a 300-mL Magnedrive II autoclave (Autoclave Engineers Inc.) equipped
FIGURE 2. Schematic diagram of experimental apparatus.
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TABLE 1. Extraction Efficiency with Variation of SCCO2 Densitya temp (°C)
pressure (atm)
density of CO2 (g/mL)
% extraction
35 35 45 45 55 55
96 129 129 170 156 201
0.7 0.8 0.7 0.8 0.7 0.8
28 28 40 51 52 57
a
[Cu2+] ) 100 ppm; HFA ) 0.74 g.
TABLE 2. Extraction Efficiency with Variation of Initial Copper(II) Concentrationa initial copper concn (ppm)
% extraction
100 200 390 650 1000
60 63 41 28 14
FIGURE 3. Path of approach to equilibrium ([Cu2+]0 ) 100 ppm; HFA0 ) 1.5 g; T ) 45 °C; P ) 136 atm). with a pressure gauge and a mixing device. All two- and three-way valves and fittings were purchased from Autoclave Engineers Inc. Aqueous copper solutions were prepared gravimetrically by dissolving copper nitrate (Aldrich Chemical Company) in deionized water. Hexafluoroacetylacetone (HFA) was obtained from Fisher Scientific Company, acetylacetone (AA) was obtained from Aldrich Chemical Company, and carbon dioxide (99%) was purchased from Connecticut Airgas Inc. For each run, 100 g of cupric nitrate solution and a known amount of chelating agent were placed into the vessel. The liquid carbon dioxide from a cylinder equipped with a dip tube was cooled in a refrigeration unit (Polar Block II, Boekel Industries, Inc.) and then charged into the vessel using a high-pressure reciprocating pump (MiniPump, Thermo Separation Products). The vessel was immersed in a water bath with an immersion circulator (Julabo) for temperature control. Temperature was monitored using a thermocouple meter (DP41-TC-MDSS, Omega Engineering) and a J-type thermocouple (Omega Engineering) inserted into a thermowell, which extended into the extraction vessel. The reactor was stirred at constant temperature and pressure for an equilibration period of 1 h, after which time the system was allowed to stand for an additional hour to permit phase separation. The separate phases were then sampled using two sampling lines: one extending deep into the reactor for sampling the aqueous phase and a shorter one used for sampling the less dense supercritical phase. The CO2 was then released, and the aqueous solution was discharged. An activated carbon bed was installed on the CO2 vent line to the fume hood. Aqueous-phase samples were analyzed directly for copper ion concentration using a Milton Roy Spectronic 601 spectrophotometer and the Bathocuproine Method, 3500-Cu E (11). The contents of the supercritical phase sample loop were first washed with ethanol, which dissolved the copper complex and the chelating agent. Ethanol was then evaporated, and the sample was mixed with a highly acidic water solution. The low pH of the solution caused reversal of the copper complex formation reaction. Nitrogen was then bubbled through the samples to remove the chelating agent before analysis by the Bathocuproine Method. Copper balance closures were better than 90%. Determination of the necessary equilibration time was done by trial. Samples were taken periodically during an extraction run. The data were then analyzed to determine the path of approach to equilibrium. Figure 3 shows the evolution of copper(II) concentration as a function of time for the copper-hexafluoroacetylacetone system at 45 °C and 136 atm. The experiment indicated that the equilibration
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a
T ) 45 °C; P ) 136 atm.
TABLE 3. Extraction Efficiency with Variiation of Initial HFA Concentrationa
a
initial HFA (g)
% extraction
1.49 0.98 0.74 0.38 0.15
60 57 44 29 18
T ) 45 °C; P ) 136 atm.
time is approximately 40 min. Therefore, all extraction runs were performed for an equilibration period of 1 h.
Discussion Results. No appreciable extraction was achieved using acetylacetone as a chelating agent. Hexafluoroacetylacetone (HFA) yielded significantly better extraction results. Subsequently, HFA was used in the experiments to determine the effects of the experimental conditions on extraction efficiency. In order to determine the effects of temperature and pressure on extraction efficiency, extractions were performed at 35, 45, and 55 °C. For each temperature, two carbon dioxide densities (pressures) were investigated, and the extraction efficiencies were determined. As shown in Table 1, at a given temperature, a change in the carbon dioxide density from 0.7 to 0.8 g/mL produced a slight rise in extraction efficiency. There is also a trend of increasing extraction efficiency with temperature for experiments conducted at the same carbon dioxide density. The effects of initial copper and HFA concentrations on extraction efficiencies were determined at 45 °C and 136 atm. Extractions were performed for a range of initial copper concentrations between 100 and 1000 ppm with 1.49 g of HFA in the system. Then, holding the initial Cu2+ concentration constant at 100 ppm, extractions were carried out for a range of initial HFA amounts between 0.15 and 1.49 g, which represent HFA excesses of approximately 200 and 2300%, respectively. The results of these experiments are given in Tables 2 and 3. Percent extraction decreased with increasing initial amount of copper and increased with increasing initial amount of HFA. Thermodynamic Modeling. Two fundamentally equivalent approaches to equilibrium calculations can be used here. Aquatic chemistry textbooks usually present an approach that
uses the law of mass action, material balances, and electroneutrality to formulate a system of as many equations as there are unknown equilibrium concentrations. The system is then solved for the molarities of all species. The alternate approach used here is modeled after the work of Walas (12) and results in a simpler system of fewer equations solvable for reaction extents. The following five aqueous-phase reactions and four phase equilibrium relations for molecular species were considered in the development of the model: Aqueous Phase Reactions:
Substituting eqs 10 and 14 into eq 13 and rearranging, the overall mole fraction becomes
∑ ν
nio +
j ij
j
Zi ) xi(β + (1 - β)Ki) )
(15)
nto +
and the aqueous-phase mole fraction of species i is given by
nio + xi )
∑ ν
j ij
j
(16)
H2O(l) a H+ + OH-
(1)
H2CO3 a HCO3- + H+
(2)
CO2(aq) + H2O(l) a H2CO3
(3)
Each mole fraction can then be converted to a molality for dilute systems using the molecular weight of water by the expression
HA(aq) a H+ + A-
(4)
mi ) 1000xi/MWH2O
2A- + Cu2+ a CuA2(aq)
(5)
Assuming an ideal system, the activity of each species in solution is equal to its molality, and for the general reaction
Phase Equilibrium Relations:
(6)
H2O* a H2O(l)
(7)
CuA2* a CuA2(aq)
(8)
HA* a HA(aq)
(9)
The asterisk (*) in the distribution relations is used to denote the species in the supercritical phase, and HA represents the chelating agent hexafluoroacetylacetone. The distribution of any species i between phases is governed by the relation:
xi* ) Kixi,aq
(10)
where Ki is the distribution coefficient and xi is the mole fraction of the ith species. The reaction extent, j, characterizes the degree of completion of a reaction at equilibrium. The number of moles of each species at equilibrium in the aqueous phase is given by
∑( ν
j i,j)
(11)
j
where nio and ni are the initial and equilibrium number of moles of the ith species, respectively. The stoichiometric number, νi,j, is positive for products and negative for reactants. Writing this expression for each species in the system and summing the moles of all species gives
nt ) nto - nt* +
(12)
where represents the algebraic sum of all the jνij terms and nto and nt are the initial and equilibrium number of moles of all species in the aqueous phase. The overall mole fraction, Zi, of any species i is given by
Zi )
nio +
ni + ni*
∑ ν
j ij
j
) nt + nt*
nto +
(13)
The phase split, β, can be defined as
β)
(17)
aA + bB T cC + dD
CO2* a CO2(aq)
ni ) nio - ni* +
(nto + )(β + (1 - β)Ki)
nt nt + nt*
(14)
the law of mass action gives
K)
aCcaDd aAaaBb
mCcmDd )
(18)
mAamBb
Substitution of the expressions for the species molalities into the law of mass action for each aqueous-phase reaction gives a set of five nonlinear equations. The only unknowns in these equations are the equilibrium extents of the five reactions. A computer program was developed to solve this system of five equations for the reaction extents using a modified Levenberg-Marquardt algorithm. The program provides the five equilibrium reaction extents, the molality of each species at equilibrium, the extraction efficiency, and pH. The initial guesses for the program were obtained using Mathemetica by solving a simplified version of the system of equations. The equilibrium constants at 25 °C for aqueous-phase reactions 1-3 and 5 were obtained from Snoeyink and Jenkins (13) and from Sekine and Ihara (8), respectively. The constant for reaction 4 was measured by measuring the pH of an aqueous solution of HFA and, lacking data to determine the standard enthalpy of formation, was assumed to remain constant with temperature. All other equilibrium constants were adjusted for temperature using standard enthalpies of formation (∆H°) by the van’t Hoff equation given by
(
)
Ki ∆H° 1 1 ln )Kref R T Tref
(19)
Since liquids are relatively incompressible, the liquid-phase equilibrium constants are weak functions of pressure, and only temperature appears in the equation. The ∆H° values for reactions 1-3 were obtained from Snoeyink and Jenkins (13) and, for reaction 5, was estimated based on documented equilibrium constant data at various temperatures (14, 15). The distribution coefficient of carbon dioxide, KCO2, was calculated using the solubility data for CO2 in water presented by Wiebe and Gaddy (16). The distribution coefficient of water, KH2O was obtained by extrapolating the solubility of water in supercritical CO2 given by Macnaughton and Foster (17). The distribution coefficient for hexafluoroacetylacetone, KHA, was not available in the literature and was therefore measured by charging the extraction vessel with water, CO2, and 1.5 g of HFA. The system was equilibrated, the phases were allowed to separate, and the supercritical phase was
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TABLE 4. Equilibrium Constants and Distribution Coefficients for Thermodynamic Model temp (°C)
pressure (atm)
log K1
log K2
log K3
log K4
log K5
KCO2
KH2O
KHA
KCuA2
35 35 45 45 55 55
96 129 129 170 156 201
-13.68 -13.68 -13.40 -13.40 -13.14 -13.14
-6.31 -6.31 -6.27 -6.27 -6.23 -6.23
0 0 0 0 0 0
-4.04 -4.04 -4.04 -4.04 -4.04 -4.04
4.00 4.00 4.14 4.14 4.28 4.28
43.74 42.29 45.44 43.22 46.35 43.89
0.0039 0.0043 0.0049 0.0056 0.0065 0.0076
0 0 0.26 0.28 0.61 0.63
65.0 92.0 116.5 140.0 179.3 224.0
TABLE 5. Sample Model Input and Output Dataa reaction H2O a + H2CO3* a H+ + HCO3H2O + CO2 a H2CO3* HA a H+ + ACu2+ + 2A- a CuA2 H+
OH-
species H+ OHH2CO3* CO2 HCO3HA ACu2+ CuA2
rxn no.
equilibrium constant
initial guess for extent
actual model extent
1 2 3 4 5
3.98 × 5.37 × 10-7 0.99949 9.12 × 10-5 1.38 × 104
2.61 × 1.38 × 10-5 0.11838 3.85 × 10-4 3.18 × 10-4
1.07 × 10-12 1.71 × 10-5 0.11831 3.54 × 10-4 2.08 × 10-4
10-14
equilibrium molality 3.71 × 1.07 × 10-11 1.18 1.18 1.71 × 10-4 5.93 × 10-2 1.46 × 10-3 5.28 × 10-4 1.54 × 10-5
10-13
species distribution coefficient
10-3
45.124 0.262 116.5
T ) 45 °C; P ) 136 atm; [Cu2+] ) 100 ppm; HFA ) 1.5 g. Extraction efficiency ) (equilibrium mol of CuA2*/initial mol of Cu2+) × 100 ) 65.4. Equilibrium pH ) -log [H+] ) 2.43. a
sampled and analyzed for HFA concentration. This procedure was carried out for the same temperatures and pressures at which the extractions were performed. No data are available in the literature relating to the partitioning behavior of the copper-chelate complex, CuA2, between carbon dioxide and water. For this reason, KCuA2 was estimated following a method developed by Brudi et al. (18), who demonstrated that the distribution coefficient of an organic compound between the SCCO2 phase and the aqueous phase can be roughly estimated by the ratio of the solubilities of the compound in the SCCO2 phase and in the aqueous phase. The solubility of the copper-HFA complex, CuA2, in SCCO2 was recently reported by Lagalante et al. (10) at 40 °C and at various pressures. The data were extrapolated to the conditions used in this study. The mole fraction solubility of CuA2 in water was measured as 3 × 10-5. Table 4 gives the values of the equilibrium constants and distribution coefficients as functions of temperature and pressure. A few other simplifications were made in the development of the model. Since the equilibrium constant for the dissociation of bicarbonate to CO32- is many orders of magnitude lower than that of the first dissociation of carbonic acid, the contribution of that reaction to overall system pH and behavior was neglected. The distribution of ionic species was also neglected in the development of the model because ionic species are insoluble in the carbon dioxide phase. Using the MINEQL equilibrium speciation model, it was determined that, at the typical pH values in the system (pH < 3), the formation of copper hydroxides and carbonates could also be neglected. The thermodynamic model successfully predicts the extraction efficiency as a function of temperature and pressure, initial amount of copper, and initial amount of HFA without the use of any adjustable parameters. Model runs at different CO2 densities yielded extraction efficiencies between 34 and 60% and are consistent with the experimental data as shown in Figure 4. The model results for the variation of the initial amount of copper and the initial amount of chelating agent are
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FIGURE 4. Comparison of experimental data and model predictions for variation of temperature and pressure ([Cu2+]0 ) 100 ppm; HFA0 ) 0.74 g). compared with the experimental results in Figures 5 and 6, respectively. A representative sample of the model output, including equilibrium concentrations, extraction efficiency, and pH, is given in Table 5. The model indicated that, if the initial amount of HFA is held constant, the equilibrium pH of the aqueous phase decreases as the initial amount of copper is increased. This behavior is a result of the formation of a greater amount of copper complex, which depletes the chelate anions and shifts the equilibrium for reaction 4 further to the right. When the model was used to simulate the effect of varying the initial amount of HFA in the system, holding initial copper concentration constant, the predicted aqueous-phase pH increased as the initial amount of HFA was decreased. The
3.07 at 45 °C and 136 atm. This value compares well with the value of 2.83 at 40 °C and 100 atm measured spectrophotometrically by Toews et al. (19). Furthermore, the model successfully predicted the inability of acetylacetone to extract copper in a single equilibrium stage even when used in large excess. For model calculations, the equilibrium constant data were obtained from Sekine and Ihara (8), and a model run was made for 35 °C and 96 atm. The computer model yielded an extraction efficiency of less than 1%, consistent with the observed phenomena. It appears that, although the formation constant of the copperAA complex is much higher than that for the copper-HFA complex, the extraction efficiency is limited by the dissociation constant of AA. Acetylacetone dissociates to a much smaller extent than HFA, providing an inadequate number of anions to complex with the aqueous copper during extraction.
Literature Cited FIGURE 5. Comparison of experimental data and model predictions for variation of initial copper concentration (45 °C and 136 atm; HFA0 ) 1.5 g).
FIGURE 6. Comparison of experimental data and model predictions for variation of initial amount of HFA (45 °C and 136 atm; [Cu2+]0 ) 100 ppm). reason for the increase in pH is that the dissociation of HFA is primarily responsible for the hydrogen ion production in the system and, with very little HFA, there are few H+ ions generated. The model proved to be fairly accurate in the prediction of system pH at equilibrium. For a simple supercritical carbon dioxide-water system, the computer model gives a pH of
(1) Lo, T. C.; Baird, M. H. I.; Hansen, C. Handbook of Solvent Extraction; Wiley Interscience: New York, 1983; p 636. (2) Ritcey, G. M.; Ashbrook, A. W. Solvent Extraction; Elsevier: Amsterdam, 1984; Part I. (3) Laintz, K. E.; Wai, C. M.; Yonker, C. R.; Smith, R. D. Anal. Chem. 1992, 64, 2875-2878. (4) Laintz, K. E.; Yu, J. J.; Wai, C. M. Anal. Chem. 1992, 64, 311-315. (5) Wang, J.; Marshall, W. D. Anal. Chem. 1994, 66, 1658-1663. (6) Lin, Y.; Wai, C. M.; Jean, F. M.; Brauer, R. D. Environ. Sci. Technol. 1994, 28, 1190-1193. (7) Laintz, K. E; Tachikawa, E. Anal. Chem. 1994, 66, 2190-2193. (8) Sekine, T.; Ihara, N. Bull. Chem. Soc. Jpn. 1971, 44, 2942-2950. (9) Cross, W.; Akgerman, A.; Erkey, C. Ind. Eng. Chem. Res. 1996, 35, 1765-1770. (10) Lagalante, A. F.; Hansen, B. N.; Bruno, T. J.; Sievers, R. E. Inorg. Chem. 1995, 34, 5781-5785. (11) Greenberg, A. E.; Clesceri, L. S.; Eaton, A. D. Standard Methods for the Examination of Water and Wastewater, 18th ed.; American Public Health Assoc., American Water Works Assoc., and Water Environment Federation: Washington, DC, 1992. (12) Walas, S. M. Phase Equilibria in Chemical Engineering; Butterworth Publishers: Boston, 1985. (13) Snoeyink, V. L.; Jenkins, D. Water Chemistry; John Wiley & Sons, Inc.: New York, 1980. (14) Serjeant, E. P.; Dempsey, B. Ionisation Constants of Organic Acids in Aqueous Solution; Pergamon Press: New York, 1979. (15) Sillen, L. G.; Martell, A. E. Stability Constants of Metal-Ion Complexes; The Chemical Society: London, 1964. (16) Wiebe, R.; Gaddy, V. L. J. Am. Chem. Soc. 1940, 62, 815-817. (17) Macnaughton, S. J.; Foster, N. R. Ind. Eng. Chem. Res. 1994, 33, 2757-2763. (18) Brudi, K.; Dahmen, N.; Schmieder, H. J. Supercrit. Fluids 1996, 9, 146-151. (19) Toews, K. L.; Shroll, R. M.; Wai, C. M.; Smart, N. G. Anal. Chem. 1995, 67, 4040-4043.
Received for review June 17, 1996. Revised manuscript received January 31, 1997. Accepted February 3, 1997.X ES960519O X
Abstract published in Advance ACS Abstracts, April 1, 1997.
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