Energy & Fuels 1997, 11, 323-326
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Ion Exchange Properties of a Western Kentucky Low-Rank Coal J. H. Kuhr,† J. D. Robertson,*,†,‡ C. J. Lafferty,‡ A. S. Wong,§ and N. D. Stalnaker§ Department of Chemistry, University of Kentucky, Lexington, Kentucky 40506-0055, Center for Applied Energy Research, University of Kentucky, Lexington, Kentucky 40511, and CST-8, MS G740, Los Alamos National Laboratory, Los Alamos, New Mexico 87545 Received September 6, 1996. Revised Manuscript Received November 18, 1996X
The ion exchange properties of a low-rank coal were evaluated to explore its viability as an inexpensive material for removing radionuclides and heavy metals from solution. Cobalt was used as a model metal to investigate the kinetics and thermodynamics of the exchange process and the effect that solution pH has on the exchange capacity. Maximum metal adsorption was found to occur at pH g3.5. The exchange was found to follow first-order kinetics and to proceed rapidly; over the temperature range investigated (-15 to 50 °C) equilibrium was reached within 10-20 min. The exchange was well described by the Langmuir adsorption isotherm with a Gibbs free energy of -8.85 ( 0.24 kJ/mol. Additional studies with environmentally significant metals and radionuclides demonstrated that cations with higher charge density were preferentially adsorbed.
Introduction Efforts to minimize production of hazardous waste and to remediate existing, accumulated waste has become one of the most important environmental challenges that the world faces today. For example, the scope and magnitude of the contamination at U.S. Department of Energy (DOE) sites are such that successful remediation of these sites will require the development of new, inexpensive materials for removing and entrapping radioactive and heavy-metal wastes from contaminated soil and water. According to a recent survey of 91 waste sites at 18 DOE facilities, the most common binary contaminant mixture reported in groundwater is a mixture of metals and radionuclides. The most frequently reported metals are lead, chromium, arsenic, and zinc, while the most commonly reported radionuclides are tritium, uranium, strontium, plutonium, cesium, cobalt, technetium, and iodine.1 The ability of low-rank coals to form stable complexes with metal ions has long been recognized.2-7 This property has been successfully utilized to estimate the concentration of acidic oxygen functional groups present in low-rank coals8 as well as serving as a convenient means of dispersing metal catalysts across a coal surface prior to liquefaction.9 Although the cation exchange capacities (CEC) of low-rank coals, typically 0.1-1 †
Department of Chemistry. Center for Applied Energy Research. Los Alamos National Laboratory. X Abstract published in Advance ACS Abstracts, January 1, 1997. (1) Riley, R. G.; Zachara, J. M. DOE/ER-0547T, 1992. (2) Eligwe, C. A.; Okolue, N. B. Fuel 1994, 73, 569-572. (3) Lafferty, C. J.; Hobday, M. D. Fuel 1990, 69, 79-82. (4) Lafferty, C. J.; Hobday, M. D. Fuel 1990, 69, 84-89. (5) Stuart, A. D Fuel 1986, 65, 1003-1005. (6) Allen, S. J. Fuel 1987, 66, 1171-1175. (7) Schafer, H. N. S. Fuel 1970, 56, 45-46. (8) Schafer, H. N. S. Fuel 1970, 49, 197-213. (9) Hatswell, M. R.; Jackson, W. R.; Larkins, F. P.; Marshall, M.; Rash, D.; Egers, E. R. Fuel 1980, 59, 442-444. ‡ §
S0887-0624(96)00146-6 CCC: $14.00
mequiv/g,3,4,10-12 are significantly lower in comparison to those of synthetic ion exchange resins, typically 7-10 mequiv/g, the substantially lower cost of the bulk material indicates great potential for the utilization of low-rank coals as a means to remove a range of metals from solution. For example, assuming the cost of the low-rank coal is $10/ton (U.S.) and that its exchange capacity is 1/10 that of a synthetic resin which costs $10/ lb, then the cost of the commercial resin system, on a dollar/mequiv basis, is 200 times that of a low-rank coalbased process. Moreover, in contrast to commercially available ion exchange resins, which produce secondary acid regeneration waste streams that require further processing, the metal-exchanged low-rank coal could be used to provide heat and power for a remediation system that would produce a nonleachable, vitrified material in which the hazardous materials are immobilized.11 In this study, the ion exchange properties of a Western Kentucky low-rank coal were evaluated to explore its viability as a remediation agent for removing heavy metals and radionuclides from aqueous solutions. Cobalt, one of the most common radionuclides encountered at DOE facilities,1 was used to investigate the chemistry of this system. The kinetic and thermodynamic properties and effect of pH and charge density on the ion exchange process of low-rank coals were investigated. In addition, the CEC of the lignite for environmentally significant metals and radionuclides was determined. (10) Kuhr, J. H.; Lafferty, C. J.; Robertson, J. D. Proceedings, Emerging Technologies in Hazardous Waste Management VII; American Chemical Society: Washington, DC, 1995; Vol. 1, pp 899-902. (11) Lafferty, C. J.; Robertson, J. D.; Kuhr, J.; Emley, J.; McGonigle, E. Proceedings, Emerging Technologies in Hazardous Waste Management VI; American Chemical Society: Washington, DC, 1994; Vol. 2, pp 807-810. (12) Robertson, J. D.; Lafferty, C. J.; Burberry, K. A. Proceedings, Emerging Technologies in Hazardous Waste Management V; American Chemical Society: Washington, DC, 1993; Vol. 2, pp 623-626.
© 1997 American Chemical Society
324 Energy & Fuels, Vol. 11, No. 2, 1997
Kuhr et al.
Experimental Section Coal Sample. The low-rank coal used as the adsorbent in these experiments was obtained from the Claiborne deposit in Carlisle County, Kentucky. A petrographical characterization of this Western Kentucky lignite can be found in the paper by Hower et al.13 When received, the as-mined coal samples were ground, passed through a 20 mesh sieve, and then stored inside sealed plastic containers to prevent moisture loss from the sample. The moisture content of the lignite ranged from 35 to 40 wt %, and the dry, ash-free oxygen content of the coal ranged from 15 to 21 wt %. No further pretreatment of the coal was performed. pH and Thermodynamic Studies. The effect of pH on the exchange capacity was investigated using solutions of 100 mg/L cobalt, as the nitrate salt, the pH of which ranged from 1 to 6. The pH of each solution was adjusted by adding either NH4OH (Baker Analyzed) or HNO3 (Baker Instra-Analyzed); the solution pH was measured using a combination singlejunction pH electrode with a Ag/AgCl reference cell and a DigiSense pH Meter (Cole-Parmer). In each case, 0.7 g of coal, dry weight, was added to 70 mL of the cobalt solution, equilibrated at 30 °C in water circulating jacketed beakers, and the resulting slurry was stirred continuously. After 2 h, the coal was removed by filtering through a 0.3 µm glass fiber filter (Gelman Sciences), and a 30 mL aliquot of the filtered solution was immediately acidified and stored in polyethylene vials for cobalt analysis. The effect of metal concentration was measured using solutions with concentrations ranging from 10 to 2000 mg/L of cobalt, which were equilibrated at 30 °C in water circulating jacketed beakers. Again, 0.7 g of coal, dry weight, was added to 70 mL of each solution, the resulting slurries were stirred continuously for 2 h, and the samples were filtered prior to analysis using a 0.3 µm glass fiber filter. As before, 30 mL aliquots of the filtered solutions for cobalt analysis were immediately acidified and stored in polyethylene vials. On the basis of the results of the pH studies, the pH values of the solutions used in the thermodynamic measurements were measured to ensure optimum metal adsorption. Because the pH values of all metal solutions used in the thermodynamic measurements exceeded 3.5, no pH adjustment was required. Reaction Kinetics. The kinetics of the exchange process were investigated using solutions of 100 mg/L of cobalt, as the nitrate salt, which were spiked with approximately 18.5 kBq of 60Co (Oxford Instruments, Inc.) and were then equilibrated at specific temperatures (-15, 0, 8, 15, 25, 35, and 50 °C) in water circulating jacketed beakers. Like the pH and thermodynamic experiments, a 1 wt % (w/v) coal slurry was used. In this case, however, 2.5 g of coal, dry weight, was added to 250 mL of solution because the measurements required that 17 1 mL aliquots be withdrawn from the slurry over the course of 1 h. As before, the solutions were stirred continuously and the pH of the solutions was adjusted to >3.5 using NH4OH to ensure optimum metal adsorption. The 1 mL aliquots were rapidly removed from contact with the coal at different time intervals using a 0.45 µm cellulose acetate syringe filter (PGC Scientifics). These samples were subsequently analyzed for cobalt by counting the radioactive tracer (60Co) remaining in solution with a NaI(Tl) well counter. The solution tested at -15 °C consisted of a 50:50 (v/v) isopropyl alcohol/water mixture that was cooled in circulating jacketed beakers with a 50:50 ethylene glycol/water mixture. To test the validity of using this mixture as a solvent system, a solution of 50:50 isopropyl alcohol/water mixture was tested at 25 °C and compared to the results of the 100% water solvent system at 25 °C. This is an important consideration since some solvents are known to cause swelling of the coal, which may in turn alter its overall CEC.3 No statistical differences were observed, at the 95% confidence level, between the two solvent systems at 25 °C. (13) Hower, J. C.; Rich, F. J.; Williams, D. A.; Bland, A. E.; Fiene, F. L. Int. J. Coal Geol. 1990, 16, 239.
Figure 1. Effect of pH on the CEC of a solution of 100 mg/L cobalt. The pH represents the pH of the cobalt solution prior to the addition of the coal. Charge Density Study. The effect of charge density on the ion exchange process was investigated using solutions of 100 mg/L of cesium, rubidium, silver, uranium, lead, cobalt, thorium, and lanthanum, all as the nitrate salts, which were prepared and equilibrated at 30 °C in water circulating jacketed beakers. The pH values of the solutions were adjusted, when necessary, to >3.5 using NH4OH. As before, 0.7 g of coal, dry weight, was added to 70 mL of each solution and the resulting slurries were stirred continuously for 2 h. The samples were then filtered using a 0.3 µm glass fiber filter, and 30 mL aliquot samples of the filtered solutions for metals analysis were immediately acidified and stored in polyethylene vials. In all of the above-mentioned studies, the metal solutions were prepared in 16 MΩ water. All experiments were conducted in the batch mode, and every reported value represents the average of at least three measurements, with the error of each measurement reported as the standard deviation of the three measurements. In every case, the samples withdrawn for analysis were immediately acidified with concentrated Baker Instra-Analyzed nitric acid to a total of 1% HNO3 (v/v). In all cases, an initial sample was taken to account for possible dilution errors introduced by pH adjustment. Also, for the kinetic studies, an initial sample was taken to correct for deviations in 60Co activity due to pipetting differences in the addition of the radioactive tracer. The metal concentration of the solutions for the pH, thermodynamic, and charge density studies were determined by flame atomic absorption and emission spectroscopy (FAA/FAES) or by direct current and inductively coupled plasma atomic emission spectroscopy (DCP/ICP-AES) as best suited by detection limit capabilities. In the case of cesium, a radioactive tracer (137Cs) and the NaI(Tl) well counter were used to measure the CEC. In the actinide studies, various amounts of the lignite (110% w/v) were stirred with the plutonium and uranium solutions for 15 min. Aliquots of these solutions were then filtered through 0.45 µm syringe filters, and the remaining actinide concentration was determined by total alpha activity using liquid scintillation counting.
Results and Discussion pH Study. As illustrated in Figure 1, the pH of the system is an important variable in the ion exchange process. The amount of cobalt adsorbed by the coal increases as the pH increases until a plateau is reached near a pH of 3.5. Previous studies on the adsorption selectivity of lignites toward a mixed solution of firstorder transition metals showed a selectivity profile increasing with atomic number, peaking at copper and then decreasing.3 This observed selectivity is consistent with the Irving-Williams order and indicates that the functional group involved in the exchange process
Ion Exchange Properties of a Low-Rank Coal
Energy & Fuels, Vol. 11, No. 2, 1997 325
Figure 2. Effect of contact time and temperature on the CEC of coal for cobalt.
contains either a nitrogen or an oxygen atom.14 FTIR studies of the raw and exchanged coals revealed that a significant amount of the metal removal observed was due to the formation of exchanged metal carboxylates on the coal surface.15 Thus, the observed dependence of the CEC of the low-rank coal on solution pH can be explained by the following reactions:
Figure 3. First-order kinetics for the initial adsorption of cobalt from solution. C0 is the initial concentration of cobalt (100 mg/L), and Ct is the cobalt concentration at time t. The -15 °C measurements were performed in a 50:50 isopropyl alcohol/water mixture. Table 1. Summary of the Rates of Reaction for the Exchange of 100 mg/L Cobalt with Coal as a Function of Temperature temp (°C) -15b
coal-COOH a coal-COO- + H+
(1)
Mn+(aq) + ncoal-COO- a (coal-COO)nM
(2)
where coal-COOH represents the carboxylic functional groups within the coal that are responsible for the metal exchange. Clearly, as the equilibrium pH of a lignite/ metal cation solution approaches or exceeds the pKa of the carboxylic acid functionalities on the coal surface, a site that is capable of forming a stable complex with the metal cations in solution is produced through the deprotonation of the carboxylic functional group. To ensure optimum metal adsorption, the kinetic, thermodynamic, and charge density experiments were performed on the plateau of the pH curve for cobalt (Figure 1). Reaction Kinetics. As indicated in Figure 2, the overall exchange process for cobalt reaches equilibrium within 10-20 min. As the temperature is increased from -15 to 50 °C, the CEC increases until it reaches a constant value at 25 °C, indicating that the energy of activation for the adsorption process has been reached. The initial rate of metal adsorption follows first-order kinetics (Figure 3), and the rate constants for the reaction at different temperatures are summarized in Table 1. The energy of activation for the adsorption of cobalt by the low-rank coal was computed by using the rate constants from first-order kinetics to construct an Arrhenius plot (Figure 4). The energy of activation for this reaction was determined to be 17.1 ( 1.1 kJ/mol, a value which is consistent with the idea that the adsorption process involves a chemical reaction.16 Note that the rate constants for 35 and 50 °C were omitted from the calculations since their corresponding rate constants are within experimental uncertainty of each other, indicating that the energy of activation for the reaction (14) Irving, H.; Williams, R. P. J. J. Chem. Soc. 1953, 3192-3210. (15) Lafferty, C. J.; Robertson, J. D.; Hower, J. C.; Verheyen, T. V. Prepr. Pap.sAm. Chem. Soc., Div. Fuel Chem. 1993, 38 (2), 468-474. (16) Ruthven, D. M. Principles of Adsorption and Adsorption Processes; Wiley: New York, 1984.
0 8 15
rate constanta (s-1) 2.95E-04 ( 0.15E-04 4.28E-04 ( 0.08E-04 5.65E-04 ( 0.25E-04 7.14E-04 ( 0.28E-04
temp (°C)
rate constanta (s-1)
25 35 50
8.19E-04 ( 0.15E-04 8.67E-04 ( 0.25E-04 8.36E-04 ( 0.21E-04
a The average and standard deviation of three or more measurements. b The solution tested at -15 °C consisted of a 50:50 (v/v) isopropyl alcohol/water mixture.
Figure 4. Arrhenius plot for cobalt adsorption.
had been exceeded at these temperatures. The error associated with the energy of activation represents the standard deviation computed from the linear regression analysis of the data presented in Figure 4. Thermodynamic Study. The effects of the metal concentration on the ion exchange properties of the coal were studied over the concentration range of 10-2000 mg/L. As illustrated in Figure 5, the adsorption data for cobalt are well described by the Langmuir adsorption isotherm model. This adsorption isotherm is given by
Ce Ce 1 ) + x/m KVm Vm where Ce is the equilibrium concentration of cobalt in solution, x is the amount of cobalt adsorbed, m is the mass of the coal, K is the monolayer binding constant, and Vm is the monolayer capacity. Using the thermodynamic relationship between the Gibbs free energy for this reaction, ∆G, and K (∆G ) -RT ln K), the Gibbs
326 Energy & Fuels, Vol. 11, No. 2, 1997
Kuhr et al. Table 3. Ability of the Western Kentucky Lignite To Remove Plutonium and Uranium from Solution % removal at a loading of solution
concn (mg/L)
initial pH
1% w/v
2% w/v
5% w/v
10% w/v
Pu3+,4+ (as nitrate) Pu3+,4+ (as acetate) Pu3+,4+ (as acetate) waste streama waste streama
880 940 5850 4.61E5b 6.05E5b
3.05 7.38 6.89 2.59 7.62
92.9 98.6 95.6
94.1 99.2 96.7
95.5 99.3 98.1 99.4 99.7
97.6 99.7 99.3
a
Waste solution of 10% HBr containing a mixture of 238Pu and The pH of the solution was increased from 2.59 to 7.62 using NH4OH. b Initial R activity of the solution in Bq/mL.
234U.
Figure 5. Langmuir adsorption isotherm for cobalt. Table 2. Summary of CEC for Various Metal Solutions with a Concentration of 100 mg/L metal cation
CECa (mequiv/g)
metal cation
CECa (mequiv/g)
Cs+ Rb+ Ag+ (UO2)2+
0.032 ( 0.008 0.037 ( 0.002 0.059 ( 0.002 0.17 ( 0.01
Pb2+ Co2+ Th4+ La3+
0.20 ( 0.01 0.24 ( 0.02 0.36 ( 0.02 0.41 ( 0.02
a The average and standard deviation of three or more measurements.
free energy for the adsorption process can be computed. The Gibbs free energy for the reaction was determined to be -8.85 ( 0.24 kJ/mol, which indicates that the adsorption process was spontaneous. Also, the fact that the data adhere to the Langmuir adsorption isotherm model, indicated by an r2 of 0.993, is consistent with a chemisorption process, a strong intermolecular interaction between the coal and cobalt.17 The error reported for the Gibbs free energy represents the standard deviation computed from the linear regression analysis of the data presented in Figure 5. Charge Density Study. The charge density of the metal cation is an important parameter in the ion exchange process as indicated in Table 2. The CEC increases as the charge density of the cation increases for all metals investigated except thorium. This reduced CEC for thorium could be due to the fact that thorium is known to form polyatomic ions in solution,18 which could reduce its charge density and, therefore, decrease its CEC. Table 2 also illustrates the applicability of this system for remediation of a range of elements such as cesium, cobalt, uranium, thorium, and lead, which are frequently the radionuclides and heavy metals of environmental concern.1 The ability of the Western Kentucky lignite to remove uranium and plutonium from solution is presented in Table 3. The results are expressed as percent removal rather than CEC as two species of plutonium (Pu3+ and Pu4+) are present in the systems investigated. As indicated by the charge density studies, the low-rank coal is very efficient in removing these highly charged cations from solution. Moreover, in contrast to the cobalt pH studies, the lignite is able to remove these actinides from solutions with a low (
