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Ind. Eng. Chem. Res. 1996, 35, 1480-1482
Use of Ozone + Hydrogen Peroxide To Degrade Macroscopic Quantities of Chelating Agents in an Aqueous Solution Evan H. Appelman,* Albert W. Jache, and John V. Muntean Chemistry Division, Argonne National Laboratory, Argonne, Illinois 60439
Ozone in the presence of hydrogen peroxide has been employed succesfully for the oxidative degradation of polycarboxylic and diphosphonic acid chelating agents present at the 0.1 M level in a neutral or weakly alkaline aqueous solution. The geminal diphosphonic acid complexants methanediphosphonic acid and (1-hydroxyethylidene)bisphosphonic acid have been successfully mineralized, as have the polycarboxylic acids ethylenediaminetetraacetic acid and tetrahydrofurantetracarboxylic acid. Introduction Chelating agents capable of complexing highly charged metal ions play an important role in separation science and in particular in the development of methods for the separation of actinide ions present in radioactive wastes. The traditional chelating agents have usually been aminopolycarboxylic acids, such as ethylenediaminetetraacetic acid, EDTA. More recently, substituted diglycolic acids, such as tetrahydrofurantetracarboxylic acid, THFTCA, have been examined (Nash et al., 1992). Also of recent interest have been geminal diphosphonic acid complexants, such as methanediphosphonic acid, MDPA, and (1-hydroxyethylidene)bisphosphonic acid, HEDPA (Nash and Rickert, 1993). These last species have an advantage over the carboxylic acid complexants in that their chelating ability is not degraded in strongly acidic solutions and also in that the free complexant retains high solubility in such solutions. All of these agents often possess a considerable degree of thermal and chemical stability, which is an advantage in their use but is an inconvenience when it is subsequently desirable to destroy the complexing ligand. Ozone in a neutral or weakly alkaline solution has been studied and used for many years as an oxidizing agent for water remediation, in some cases as an alternative to chlorination of water (Horvath et al., 1985). More recently, improved variations of the technology have been suggested and tested. The introduction of hydrogen peroxide as a coreagent leads to the formation of OH radicals (Staehelin and Hoigne´, 1982; Tomiyasu et al., 1985), and in water remediation the combination of ozone and hydrogen peroxide has been found to be considerably more effective than ozone alone for the removal of low levels of organic material (Glaze et al., 1987; Aieta et al., 1988; Glaze and Kang, 1989). Ultraviolet irradiation of aqueous ozone produces hydrogen peroxide in situ (Taube, 1957) and also results in enhanced destruction of organic matter (Peyton et al., 1982; Glaze et al., 1987). Finally, TiO2 has been used as a heterogeneous catalyst for oxidation of aqueous organic matter by ozone both with and without UV irradiation (Ollis, 1985; Allemane et al., 1993). All of these methods have generally been concerned with the oxidation of “micropollutants”, i.e., organic compounds present at millimolar concentrations or below. To the best of our knowledge, ozone has not heretofore been shown to be useful for the oxidation of substantial quantities of organic material in water, i.e., at concentrations of 0.1 M or greater. In particular, no investigations have been carried out of the efficacy of this oxidant for the destruction of macroscopic concen0888-5885/96/2635-1480$12.00/0
trations of aqueous chelating agents. In the present paper, we report the use of the combined ozonehydrogen peroxide system for oxidative degradation of 0.1 M aqueous solutions of the chelating agents HEDPA, MDPA, THFTCA, and EDTA. Experimental Section HEDPA was obtained from Albright and Wilson Americas, Inc., as a 65 wt % aqueous solution and was purified by crystallization from acetic acid. The other chelating agents studied were purchased from Aldrich and used without further purification. Ozone was prepared by passing high-purity oxygen through a homemade ozone generator of the conventional silentdischarge type. The incoming oxygen flow was measured with a thermal-conductance mass flowmeter, and ozone output was determined by bubbling the effluent gas for a measured time into a neutral KI solution, followed by acidification and titration of the liberated iodine with standard thiosulfate (Kolthoff and Belcher, 1957). An ozone concentration of approximately 5 mol % was routinely achieved, corresponding to an output of about 0.3 mmol of ozone/min. Hydrogen peroxide was determined by its molybdatecatalyzed reaction with excess iodide in an acid solution, followed by titration of the liberated iodine with a standard thiosulfate solution (Kolthoff and Belcher, 1957). The phosphate present at the end of an HEDPA or MDPA reaction was determined spectrophotometrically as the phosphomolybdovanadate complex (Quinlan and DeSesa, 1955). (This spectrophotometric procedure could only be used after complete reaction and decomposition of the hydrogen peroxide, since otherwise the peroxovanadium complexes would interfere with the spectrophotometric analysis.) The course of the destruction of organic substrates was followed by 1H and, for the diphosphonates, by 31P NMR spectroscopy. Samples were made ca. 20% in D2O to provide a lock for the spectrometer. In the case of HEDPA and MDPA, the 31P spectra were depended on for quantitation, and the proton spectra were only used for qualitative and semiquantitative estimation of intermediate products. For THFTCA and EDTA reactions, the proton spectra were quantitated by comparison with an added standard. NMR measurements were carried out with a GE gn-based Model 293 Omega spectrometer (7.06 T) at Larmor frequencies of 300.52 and 121.65 MHz for protons and 31P, respectively. Reactions were carried out by passing the ozone/ oxygen mixture through a thin glass capillary into a substrate solution held in a constant-temperature water © 1996 American Chemical Society
Ind. Eng. Chem. Res., Vol. 35, No. 4, 1996 1481 Table 1. Oxidation of HEDPA and MDPA by Ozone + Hydrogen Peroxidea
Table 2. Reaction of EDTA and THFTCA with Ozone + Hydrogen Peroxidea
pH [H2O2]0 residual reaction O3 utilization (%)c (M) initial final substrate (%) time (min)b
pH [H2O2]0 reaction O3 utilization residual (%)d (M) initial final organic (%)b time (min)c
0.6d 0.6d 0d 0.05d 0.1 0.2
∼11.6e ∼11.6 9.7g 9.3 9.5 8.2 9.5 7.4 9.3 6.8 9.1 6.7
0.2 0.4 0.8 1.7
9.1 9.2 9.1 9.1
7.5 7.4 7.3 7.7
HEDPA 78f 56h 88 44h 14h 0i
6 28 20 30 67 70
15 6.3 2.4 11.2 7.7 8.6
0.25 0.5 0g 0.25 0.5 0.25
MDPA 45 27 14 5
33 43 55 71
6.7 6.8 6.3 5.4
1.0g 1.2g 0.25 0.5 0g 0.133 0.25 0.5 0.25 0.5
a Unless otherwise specified, reactions consisted of treatment of 3 cm3 of a 0.1 M solution in substrate at 25 °C with 150 cm3/ min of 5 mol % O3 in O2 (ca. 0.3 mmol of O3/min). Unless use of a buffer is designated, aqueous NaOH was added to adjust the initial pH of the solution. b Until H2O2 was completely consumed; a comparable time period was used in the absence of H2O2. c With respect to the amount of ozone delivered into the reaction mixture. Assumes two oxidizing equivalents per mole of O3 and oxidation of substrate to phosphate, CO2, and water. d Reaction volume of 2 cm3 at ca. 23 °C. e Buffered by 0.5 M H2O2/0.5 M HO2-. f Four minor unidentified 31P NMR resonances were seen, corresponding to a total of ca. 10% of the initial HEDPA, along with small amounts of acetate and unidentified species seen in the proton spectrum. g Buffered by 0.08 M borax. h Small amounts of formate and acetate seen in the proton NMR spectrum. i Small amounts of acetate seen in the proton NMR spectrum.
9.1e 9.1e 7.0h 6.9h 6.9h 6.9i
9.0 9.0 7.2 8.2 8.5 7.4
∼11.6j ∼11.6 ∼8.5l ∼9.6 8.7e 9.0 8.7e 9.1 7.5n 7.5 7.5n 7.8 7.4n 8.0 7.3n 8.1 h 6.8 8.8 6.8h 9.1
EDTA 9f 6f 14f 5f 0 10f
40 40 45 37 80 105
48 48 28 53 25 18
THFTCA ∼80k 22m 24m 24m ∼100o 14m 3m 0 14m 4p
7 ∼30 25 25 72 160 152 154 60 80
21 22 39 39 0 6.4 7.1 7.1 17 13.5
a Unless otherwise specified, reactions consisted of treatment of 3 cm3 of a 0.1 M solution in substrate at 25 °C with 150 cm3/ min of 5 mol % O3 in O2 (ca. 0.3 mmol of O3/min). b Based on the sum of integrated proton NMR resonances. Except as noted, the initial substrate was entirely consumed. c Until H2O2 was completely consumed; a comparable time period was used in the absence of H2O2. d With respect to the amount of ozone delivered into the reaction mixture. Assumes two oxidizing equivalents per mole of O3 and oxidation of substrate to CO2, N2, and water. Also assumes that unidentified intermediate degradation products have on average half the reducing capacity of the original substrate. e Self-buffered. Aqueous NaOH was added to adjust the initial pH of the solution. f Mixture of degradation products, including acetate and formate. g Reaction volume 2 cm3 at ca. 23 °C. h Buffered by 0.1 M H PO -/0.1 M HPO 2-. i Buffered by 0.3 M 2 4 4 H2PO4-/0.3 M HPO42-. j Buffered by 0.5 M H2O2/0.5 M HO2-. k Unreacted THFTCA plus complex product mixture. l Buffered by 0.2 M potassium tetraborate. m Complex mixture of degradation products, including malonate, acetate, and formate. n Buffered by 0.4 M H2PO4-/0.8 M HPO42-. o Unreacted starting material. p Two dominant products, one of which is malonate (see text).
bath. Samples were extracted at periodic intervals for determination of the hydrogen peroxide concentration, and reaction was continued until the hydrogen peroxide had been completely or very nearly completely consumed, at which time the product solutions were characterized by NMR spectroscopy, and those resulting from oxidation of HEDPA and MDPA were analyzed spectrophotometrically for phosphate. In a few experiments NMR spectra were measured during the course of the reactions, in tandem with the iodometric titrations.
presence of substrate, presumably as the result of scavenging of radical intermediates.
Results
Discussion
The results for the phosphonate and carboxylate complexants are summarized in Tables 1 and 2, respectively. The experiments in which NMR spectra were measured during the course of the reactions indicated the consumption of hydrogen peroxide to be at least a qualitative monitor of the progress of oxidation of the substrate. In the case of MDPA neither 1H nor 31P NMR spectra showed significant concentrations of any organic or phosphorus-containing species in solution other than starting material and phosphate. In the case of HEDPA, minor amounts of formate and acetate were formed, particularly in those reactions in which not all of the starting material had been consumed. When the oxidation of HEDPA was carried out at very high pH, four new phosphorus resonances appeared, indicating the formation of phosphorus-containing intermediate oxidation products. In the case of EDTA and THFTCA, on the other hand, the starting material was often completely consumed, and the residual organic material appeared to be distributed among a number of intermediate oxidation products, only a few of which were we able to characterize. We observed in general that the rate of disappearance of hydrogen peroxide was considerably decreased by the
It is evident from our results that it is indeed possible to mineralize substantial concentrations of both carboxylic and phosphonic acid complexants through the use of a mixture of hydrogen peroxide and ozone. The beneficial effect of the peroxide is evident from the data. Even in the case of EDTA, which is destroyed by ozone alone, satisfactory cleanup of intermediate oxidation products requires the presence of H2O2. The oxidation of MDPA appears to require particularly forcing conditions, nearly complete oxidation of the substrate being only attainable in the presence of around 2 M hydrogen peroxide. The completeness of oxidation of substrates tends to decrease with increasing pH, presumably because of the scavenging of OH radicals by HO2- and CO32-. At low pH, on the other hand, the overall reaction becomes slower. Hence, in any practical application, it will be necessary to choose an optimal pH at which to carry out the reaction. The oxidation of “macropollutants” results in substantial pH changes as reaction proceeds. Oxidation of phosphonates tends to drive the pH to lower values, because of the replacement of the phosphonate ions by more acidic phosphate ions, while oxidation of carboxylates increases the pH, because of the continual flushing
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of carbon dioxide from the solution. Self-buffering or the use of external buffers is a highly imperfect means of pH control, and in a practical system it will be desirable to maintain an optimal pH by continual addition of acid or base. The utilization of ozone is often relatively low in our experiments. We have not, however, optimized ozone delivery, and it may be possible to improve the efficiency of ozone utilization by employing a gas dispersal system that more effectively keeps the solution saturated in ozone. In most of the systems, the utilization of ozone decreases as the initial H2O2 concentration increases. This is not surprising, since both H2O2 and HO2- react with OH radicals. Hence, it should be advantageous to maintain a relatively low hydrogen peroxide concentration by gradual addition of H2O2 as reaction proceeds. The simultaneous addition of H2O2 and ozone has already been proposed for the use of this oxidant system in water remediation (Glaze and Kang, 1989). In the real world, complexants are likely to be accompanied by significant concentrations of transitionmetal ions, which could have detrimental effects on the oxidation process. We have not addressed this problem in the present preliminary study, but it needs to be considered in further investigations. Conclusions The combination of ozone with hydrogen peroxide is already known to be one of several techniques that enhance the action of ozone itself in the destruction of low levels of organic compounds in water. The results of the present preliminary study suggest that it should be practical to apply the method to the destruction of macroscopic quantities (0.1 M or greater) of aqueous organic materials and in particular to the mineralization of carboxylic and phosphonic acid complexants in an aqueous solution. Acknowledgment We thank Dr. K. L. Nash for suggesting the use of the ozone-hydrogen peroxide oxidant system to destroy substantial concentrations of aqueous complexants and Dr. Dan Meisel for a very helpful critique of the manuscript. This research was performed under the auspices of the Office of Basic Energy Sciences, Division of Chemical Sciences, U.S. Department of Energy, under Contract W-31-109-Eng-38.
Literature Cited Aieta, E. M.; Reagan, K. M.; Lang, J. S.; McReynolds, L.; Kang, J.-W.; Glaze, W. H. Advanced Oxidation Processes for Treating Groundwater Contaminated with TCE and PCE: Pilot-Scale Evaluations. J. Am. Water Works Assoc. 1988, 80, 64-72. Allemane, H.; Delouane, B.; Paillard, H.; Legube, B. Comparative Efficiency of Three Systems (O3, O3/H2O2 and O3/TiO2) for the Oxidation of Natural Organic Matter in Water. Ozone: Sci. Eng. 1993, 15, 419-432. Duguet, J. P.; Brodard, E.; Dussert, B.; Mallevialle, J. Improvement in the Effectiveness of Ozonation of Drinking Water Through the Use of Hydrogen Peroxide. Ozone: Sci. Eng. 1985, 7, 241-258. Glaze, W. H.; Kang, J.-W. Advanced Oxidation Processes. Description of a Kinetic Model for the Oxidation of Hazardous Materials in Aqueous Media with Ozone and Hydrogen Peroxide in a Semibatch Reactor. Ind. Eng. Chem. Res. 1989, 28, 1573-1587. Glaze, W. H.; Kang, J.-W.; Chapin, D. H. The Chemistry of Water Treatment Processes Involving Ozone, Hydrogen Peroxide and Ultraviolet Radiation. Ozone: Sci. Eng. 1987, 9, 335-352. Horvath, M.; Bilitzky, L.; Hu¨ttner, J. Ozone; Elsevier: Amsterdam, The Netherlands, 1985. Kolthoff, I. M.; Belcher, R. Volumetric Analysis; Interscience: New York, 1957; Vol. III, pp 281-283. Nash, K. L.; Rickert, P. G. Thermally Unstable Complexants: Stability of Lanthanide/Actinide Complexes, Thermal Instability of the Ligands, and Applications in Actinide Separations. Sep. Sci. Technol. 1993, 28, 25-41. Nash, K. L.; Horwitz, E. P.; Gatrone, R. C.; Rickert, P. G. The Effect of Ligand Rigidity on the Stability of Europium(III) Complexes with Substituted Diglycolic Acids. J. Alloys Compd. 1992, 180, 375-381. Ollis, D. E. Contaminant Degradation in WatersHeterogeneous Photocatalysis Degrades Halogenated Hydrocarbon Contaminants. Environ. Sci. Technol. 1985, 19, 480-484. Peyton, G. R.; Huang, F. Y.; Burleson, J. L.; Glaze, W. H. Destruction of Pollutants in Water with Ozone in Combination with Ultraviolet Radiation. 1. General Principles and Oxidation of Tetrachloroethylene. Environ. Sci. Technol. 1982, 16, 448453. Quinlan, K. P.; DeSesa, M. A. Spectrophotometric Determination of Phosphorus as Molybdovanadophosphoric Acid. Anal. Chem. 1955, 27, 1626-1629. Staehelin, J.; Hoigne´, J. Decomposition of Ozone in Water: Rate of Initiation by Hydroxide Ions and Hydrogen Peroxide. Environ. Sci. Technol. 1982, 16, 676-681. Taube, H. Photochemical Reactions of Ozone in Solution. Trans. Faraday Soc. 1957, 53, 656-665. Tomiyasu, H.; Fukutomi, H.; Gordon, G. Kinetics and Mechanism of Ozone Decomposition in Basic Aqueous Solution. Inorg. Chem. 1985, 24, 2962-2966.
Received for review June 19, 1995 Revised manuscript received October 19, 1995 Accepted December 19, 1995X IE950367Q X Abstract published in Advance ACS Abstracts, February 1, 1996.
