Investigation of Mechanisms of Oxidation of EDTA and NTA by

Jul 13, 2006 - Department of Civil and Environmental Engineering,. Box 352700, University of Washington,. Seattle, Washington 98115-2700. Permanganate...
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Environ. Sci. Technol. 2006, 40, 5089-5094

Investigation of Mechanisms of Oxidation of EDTA and NTA by Permanganate at High pH HYUN-SHIK CHANG, GREGORY V. KORSHIN,* AND JOHN F. FERGUSON Department of Civil and Environmental Engineering, Box 352700, University of Washington, Seattle, Washington 98115-2700

Permanganate has been used for oxidation of nuclear wastes containing chelating agents such as ethylenediaminetetraacetic and nitrilotriacetic acids (EDTA and NTA) to improve separation of radionuclides and heavy metals from the wastes, but the mechanisms of degradation of these and related organic ligands at high pHs have not been studied. EDTA, NTA, and the model compound ethylenediamine (EN) were found to be readily oxidized by permanganate at pH 12-14. The reduction of permangante was accompanied by formation of unstable manganate and dispersed MnO2 particles, which constituted the final product of permanganate reduction. The yields and speciation of EDTA, NTA, and EN breakdown products were affected by the pH and permanganate dose. Iminodiacetic acid (IDA), oxalate, formate, and ammonia were the predominant EDTA and NTA oxidation products. Mineralization of EDTA, NTA, and EN to CO2 was more significant at pH 12. At pH 14 formation of oxalate and deamination to NH3 were the most important reactions. IDA was released upon the oxidation of both EDTA and NTA, but EDTA oxidation yielded no ethylenediaminediacetic acid (EDDA). The speciation of the reaction products indicated that the ethylene group in EDTA was the preferred attack site in oxidations by alkaline permanganate.

ganese solids that bind TRU and some of the other target species (1-4), but intrinsic mechanisms of these processes have not been explored. In acidic or circumneutral solutions, the oxidation of EDTA is accompanied by release of ethylenediamine-N,N′,N′-triacetic acid (ED3A) and CO2 as the major products of EDTA degradation, and Mn2+ as the final product of Mn(VII) reduction (8). Microbiologically mediated degradation of EDTA in circumneutral pHs also proceeds via the formation of ED3A, N,N′-ethylenediaminediacetic acid (EDDA), iminodiacetatic acid (IDA), and low-molecularweight compounds (9, 10). Experiments with model compounds such as amino acids (e.g., 11-14) have shown that the reduction of permanganate by these species at high pHs proceeds via formation of manganate MnO42-, hypomanganate MnO43- (15), and finally MnO2 or mixed Mn(III)/Mn(IV) solids (16). Identified reaction products formed upon the permanganate oxidation of amino acids and other model compounds include ammonia, oxalate, CO2, and traces of aldehydes (17-20). Similar species and, in addition to them, IDA, glycine, and glycolate were identified in solutions of EDTA oxidized at high pHs by Ag(III) (21), and also at circumneutral pHs in the case of oxidation by hydroxyl radicals generated in the H2O2/UV system (22) or with reactive oxygen species activated by zerovalent iron (23). Despite the extent of studies concerned with the microbiological and chemical degradation of EDTA, NTA, and other persistent organic ligands at acidic and circumneutral conditions, the nature of processes that govern the oxidative breakdown of such compounds at high pHs has not been ascertained, nor have the identities and yields of breakdown products in these conditions been quantified. The goal of this study was to explore in detail the degradation of EDTA, NTA, and a model compound ethylenediamine (EN) in alkaline media. In addition to the practical utility of such data for modeling of physicochemical processes in nuclear wastes and compartments of the environment affected by them (e.g., the subsurface zone), more detailed and chemically explicit theories of EDTA, NTA, and EN oxidation at high pHs will enhance and complement the current understanding of the processes that govern the environmental fate of persistent organic ligands.

Experimental Section Introduction Ethylenediaminetetraacetic and nitrilotriacetic acids (EDTA and NTA, respectively) are found at concentrations as high as 0.03 M in many tanks containing nuclear wastes located at the Hanford Nuclear Reservation (1-4). Formation of complexes of EDTA, NTA, and related ligands with transuranium elements (TRU), 90Sr, other radionuclides, and metals (for instance, chromium, nickel, copper) interferes with removal of the target species from the wastes, which is necessary to facilitate waste handling, maximize vitrification of radionuclides for long-term disposal, and improve the long-term stability of the vitrification products. These complexes also tend to persist in the environment and exhibit a greater mobility in surface waters and the subsurface zone (5-7). Permanganate treatment of the high pH, high ionic strength wastes has been shown to improve separation of the target species because of the breakdown of EDTA (and other organic ligands present in the wastes) caused by permanganate oxidation and attendant formation of man* Corresponding author e-mail [email protected]; phone: (206) 543-2394; fax: (206) 685-9185. 10.1021/es0605366 CCC: $33.50 Published on Web 07/13/2006

 2006 American Chemical Society

All chemicals were ACS reagent grade. Reagent water was obtained from a Millipore Super-Q water system. All experiments were conducted at 20 °C and a constant ionic strength (1.13 M) which was adjusted with sodium perchlorate. Concentrations of sodium hydroxide in solutions with pH values of 12, 13, and 14 were 0.01, 0.1, and 1 M, respectively. EDTA, NTA, IDA, and EN oxidation experiments concerned with the quantification of reaction product formation were carried out using 9 mL disposable test tubes to which varying amounts of permanganate were added. Following that, the test tubes were mildly agitated for 40 h to complete reactions between permanganate and organic substrates. Manganese solids formed during the reaction were allowed to precipitate. Changes of pH values during the reaction were less than 0.2 pH units. The supernatant was filtered through a 0.45 µm filter. The retained solids were used in selected cases for morphological and structural analyses. Concentrations of EDTA, NTA, ethylenediaminediacetic acid (EDDA), IDA, acetate, glycolate, formate, oxalate, nitrate, and nitrite were measured using a Dionex DX500 ion chromatograph equipped with a CD20 conductivity detector, GP40 gradient pump, and IonPac AS11 separation column. The loop size was 24 µL. Three degassed eluents, including VOL. 40, NO. 16, 2006 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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FIGURE 1. Absorbance spectra of EDTA/permanganate system for the first 60 min of reaction and pH 12 (A) and 14 (B). EDTA concentration 10-4 M, permanganate/EDTA molar ratio 1.0. 191 mM NaOH (A), DI water (B), and 19 mM (C) NaOH solutions, were mixed in the gradient mode. The initial set up was A 0%, B 98.5%, and C 1.5%, and it was ramped to A 30%, B 70%, C 0% with time. Flow rate was set as 0.3 mL/ min, and total run time was 75 min. Prior to analyses, the samples were all filtered with a 0.45 µm filter and passed through OnGuard II H cartridges to lower their pH. The concentrations of EN were estimated based on the absorbance of its complex with copper at wavelength 546 nm. Total organic carbon (TOC) was measured using an OI model 700 TOC Analyzer. Mineralization of organic substrates to CO2 was quantified as the difference between TOC concentration in solutions of EDTA, NTA, and EN before and after permanganate treatment. Concentration of ammonia was measured using Standard Method 10031 with a HACH DR/ 4000 spectrophotometer. Absorbance spectra were acquired using a Perkin-Elmer Lambda-18 UV/VIS spectrophotometer. Concentration of soluble manganese was measured with a Jobin-Yvon JY-138 Ultrace inductively coupled plasma emission spectrometer.

Results Interactions between EDTA, NTA, EN, and permanganate were in all cases accompanied by the reduction of the permanganate, attendant degradation of the target compounds, and formation of a variety of reaction products. Manganese species formed upon the reduction of permanganate could be readily identified by UV-Vis spectrophotometry due to their distinct absorbance spectra and high molar extinction coefficients (24). The absorbance spectra of the EDTA/MnO4- system showed a rapid consumption of permanganate in the first hour of the reaction. The reduction of permanganate could be quantified by the disappearance of the distinct permanganate absorbance band located in the range of wavelengths 490 to 590 nm and formation of bands with maxima at 347, 439, and 606 nm characteristic of manganate MnO42- as shown in Figure 1A and B for pHs 5090

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FIGURE 2. Effect of permanganate/ethylenediamine molar ratio on molar yields of oxalate and NH3 at varying pHs. Initial EN concentration 10-2 M, time 40 h. 12 and 14, respectively. The apparent rate of the reduction of permanganate to manganate increased with the pH. The stability of the manganate released during the initial phase of the reaction also increased at higher pHs, as evidenced by the nearly constant intensity of the manganate absorption bands at pH 14 during the first 60 min of reaction (Figure 1B) and their decay at pH 12 (Figure 1A) that is indicated by the behavior of the MnO42- band at 606 nm. No evidence for hypomanganate MnO43- with its maxima located at 313 and 667 nm was seen in these experiments. Absorbance measurements for reaction times between 1 and 24 h showed that the concentrations of permanganate and manganate gradually decreased at all pHs and permanganate doses. For a 40 h reaction time, permanganate and manganate were completely consumed and the concentration of soluble manganese was below its detection limit ( 4 (Figure 6). The disappearance of EDTA was accompanied by the release of CO2, IDA, oxalate, formate, glycolate, nitrite, nitrate, and ammonia. The total concentration of organic carbon gradually decreased for all MnO4-/EDTA molar doses. The molar yields of CO2 decreased with the pH and at the highest permanganate dose they were 5.0, 3.4, and 1.9 for pH 12, 13, and 14, respectively (Figure 6C). In contrast with that, the yields of oxalate increased with the pH to reach values of 1.1, 2.3, and 3.0 for pH 12, 13, and 14, respectively (Figure 6B). The release of formate was notable among minor products for MnO4-/EDTA molar ratios < 2, with yields of formate slightly less than 0.1 at pH 12 for a 2.0 MnO4-/EDTA molar ratio (Figure S4). The concentration of formate declined VOL. 40, NO. 16, 2006 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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FIGURE 5. Nitrogen (A) and carbon (B) mass balances for NTA at pH 14 and varying MnO4-/NTA molar ratios. rapidly at higher permanganate doses and pHs. The yield of glycolate was much lower than that of formate having a maximum of 0.04 at pH 14 and a 0.5 MnO4-/EDTA molar ratio. Nitrogen-containing species formed upon the oxidation of EDTA by permanganate were predominated by IDA and ammonia with lower levels of nitrate (Figure 7A). Only traces of nitrite were found. Calculations of the carbon and nitrogen balances for EDTA degraded by permanganate (Figure 7B and Figure S5) indicated that, as was the case with NTA, IDA, oxalate, and formate were the main identified intermediate species. The contribution of unidentified species was noticeable for MnO4-/EDTA molar ratios < 4 and pH < 14. We hypothesized that ED3A was a major contributor to the unidentified fraction of the carbon and nitrogen shown in Figure 7. The formation of ED3A via permanganate oxidation of EDTA in acidic and circumneutral pHs was noted in prior research (e.g., 8, 9, 22); however, the occurrence of this compound was not quantified in this study. For the highest MnO4-/EDTA molar ratio of 8 and pH 14, oxalate, ammonia, and carbon dioxide contributed to the almost complete carbon and nitrogen mass balance in the system.

Discussion Comparison of the data for EN, NTA, and EDTA shows that at highly basic conditions, the degradation of EDTA by permanganate is likely to proceed via N-dealkylation that starts at the ethylene group between the nitrogen atoms, as was observed with UV/H2O2 in the absence of iron and with oxygen activation by Fe(0) (22, 23). However, side-chain N-dealkylation to ED3A and EDDA observed with permanganate at circumneutral pH (8, 10, 22, 23) cannot be ruled out, since mass balances on C and N are incomplete at low MnO4-/EDTA ratios. At higher ratios and pH 14, mass balances are complete and support the oxidation of the ethylene group. The oxidation and deamination of EN also 5092

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FIGURE 6. Effect of permanganate/EDTA molar ratio and pH on the degradation of EDTA and formation of oxalate and carbon dioxide. Initial EDTA concentration 10-2 M, reaction time 40 h. demonstrates the occurrence of such a reaction. Mechanistically, this reaction is likely to follow the classical pathway of sequential one-electron-transfer steps that involve the formation of transient imines, their deprotonation, and hydrolysis (25-27). A possible sequence of these steps for EN is shown in Figure S6 in the Supporting Information. Further oxidation of the carbon in glyoxal formed as a transient product during N-dealkylation/oxidation of EN is also shown in Figure S6. Glyoxal or glycolate that was detected in small amounts in experiments with EN, NTA, and EDTA can be further oxidized to formate, oxalic acid, and CO2. It can also be hypothesized that the predominance of oxalate (rather than formate) release for permanganatetreated EN indicates that the first one-electron transfer that initiates the oxidation entire sequence is not necessarily followed by a second electron transfer at the same site but rather by the oxidation of the N′-nitrogen atom and formation of a transient N,N′-diimine that breaks down to release ammonia and precursors of oxalate (for instance, glyoxal). The formation of a transient N,N′-diimine is denoted as pathway B in Figure S6. The pathways of dealkylation differed somewhat between NTA and EDTA. The oxidation of NTA in which IDA was the

This hypothesis does not preclude the formation of ED3A and, less likely, EDDA and their ensuing oxidation to IDA and other byproducts. That pathway is likely to become increasingly important at pH < 14, as indicated by a comparatively larger fraction of unidentified products possibly associated with ED3A at pH 12 (Figures S4 and S5) and high ED3A yields at circumneutral pHs (8). However, the notable absence of EDDA and presence of IDA at pH 14 indicate that the preferred locus of EDTA oxidation at high pHs is the central ethylene group. A possible reaction sequence for the formation of IDA upon the oxidation of EDTA by permanganate is shown in Figure 8. In this figure, pathways A and B denote the breakdown of EDTA via electron transfer steps that occur sequentially either at the same nitrogen atom, or via the formation of a transient N.N′-diimine structurally similar to that hypothesized for EN (Figure S6). An alternative reaction sequence that includes the apparently less preferred pathway of EDTA oxidation via the formation of ED3A is shown in Figure S7. MnO4- is expected to be the acceptor of electrons but not necessarily the actual oxidizing agent in the reaction sequences shown in Figure 8 and Figures S6 and S7. The reduction of MnO4- results in formation of manganate that is evident from the absorbance spectra shown in Figure 1. The manganate then disproportionates to form permanganate and manganese dioxide (28), as shown in eq 1.

3MnO42- + 2H2O f 2MnO4- + MnO2 + 4OH- (1) FIGURE 7. Nitrogen (A) and carbon (B) mass balances for EDTA at pH 14 and varying MnO4-/EDTA molar ratios. most prominent intermediate product (Figure 5) clearly followed the classical pathway of N-dealkylation. For EDTA, side-chain N-dealkylation would have yielded ED3A and then EDDA as its typical oxidation products. However, no EDDA was found in solutions of EDTA treated with permanganate. The occurrence of relatively high concentrations of IDA and, at the same time, the absence of EDDA, showed that at pH 12-14 the attack on the ethylene group is more prevalent than the N-dealkylation of the side acetate groups in EDTA.

The reaction scheme shown in Figure 8 is similar to that developed to explain the mechanisms of photochemical degradation of EDTA, in which hydroxyl radical was the main oxidizing agent and IDA was the most prominent breakdown product (22). The generation of hydroxyl radicals preceded by formation of a transient monohydroxo complex of permanganate and followed by the release of manganate (eqs 2 and 3) was hypothesized to occur in oxidations by alkaline permanganate (19, 29), in agreement with earlier results that demonstrated that at high pHs almost all of the oxygen transferred to the organic substrate being oxidized originated from water molecules rather than from the

FIGURE 8. Suggested pathway of permanganate EDTA oxidation and IDA formation at high pHs via oxidative attack on the ethylene group. VOL. 40, NO. 16, 2006 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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permanganate per se (28).

MnO4- + OH - T [MnO4‚OH]2[MnO4‚OH]

2-

f MnO4

2-

+ ‚OH

(2) (3)

In the case of permanganate oxidation of EDTA and NTA in the Hanford nuclear wastes, the rapid breakdown of the chelating agents and attendant formation of MnO2 solids will drive both the decomplexation of the target metal ions and their adsorption on the surfaces of the manganese solids. It remains to be determined whether complexation of the target cations with IDA or other possibly present degradation products (EDDA, ED3A) can interfere with their removal by adsorption, or whether the formation of EDTA complexes with the relevant metal cations (for instance, Ni2+, Cu2+, Cr3+) found in the Hanford nuclear affects the kinetics of their oxidation and the nature of the breakdown products.

Acknowledgments This study was supported by Consortium for Risk Evaluation and Stakeholder Participation (CRESP), Department of Energy Grant DE FG01-03EW1536. We thank Professor James Mayer (University of Washington) and Dr. Andy Felmy (Pacific Northwest National Laboratory) for their participation in the discussion concerning the nature and significance of oxidative processes for Hanford nuclear wastes’ treatment.

Supporting Information Available Six additional figures demonstrating mineralization of EN at varying pHs, carbon and nitrogen balances for EN and EDTA at pHs 12 and 13, and a hypothetical pathway of degradation of EN by permanganate; a figure representing an alternative mechanism of EDTA degradation via the formation of ED3A. This material is available free of charge via the Internet at http://pubs.acs.org.

Literature Cited (1) Felmy, A. R.; Qafoku, O. An aqueous thermodynamic model for the complexation of nickel with EDTA valid to high base concentrations. J. Solution Chem. 2004, 33 (9), 1161-1180. (2) Felmy, A. R.; Mason, M. J. An aqueous thermodynamic model for the complexation of sodium and strontium with organic chelates valid to high ionic strengths. I. Ethylenedinitrilotetracetic acid (EDTA). J. Solution Chem. 2003, 32 (4), 283-300. (3) Bond, A. H.; Nash, K. L.; Gelis, A. V.; Sullivan, J. C.; Jensen, M. P.; Rao, L. Plutonium mobilization and matrix dissolution during experimental sludge washing of bismuth phosphate, redox and purex waste simulants. Sep. Sci. Technol. 2001, 36, (5/6), 12411256. (4) Hallen, R. T.; Geeting, J. G. H.; Lilga, M. A.; Hart, T. R.; Hoopes, F. V. Assessment of the mechanisms for Sr-90 and TRU removal from complexant-containing tank wastes at Hanford. Sep. Sci. Technol. 2005, 40 (1), 171-183. (5) Nowack, B.; Van Briesen, J. M. Chelating agents in the environment. In Biogeochemistry of Chelating Agents; ACS Symposium Series 910; American Chemical Society: Washington, DC, 2005; pp 1-19. (6) Mayes, M. A.; Yin, X. L.; Pace, M. N.; Jardine, P. M. Rates and mechanisms of Co(II)EDTA2- interactions with sediments from the Hanford site. In Biogeochemistry of Chelating Agents; ACS Symposium Series 91; American Chemical Societ: Washington, DC, 2005; pp 278-296. (7) Mayes, M. A.; Mehlhorn, T. L.; Jardine, P. M. Coupled hydrological and geochemical processes influencing the transport of chelated metals in the ORNL vadose zone and groundwater. In Biogeochemistry of Chelating Agents; ACS Symposium Series 910; American Chemical Society: Washington, DC, 200; pp 297315. (8) Bose, R. N.; Keane, C.; Xidis, A.; Reed, J. W.; Li, R. M.; Tu, H.; Hamlet, P. L. Oxidation of etylenediaminetetraacetic acid by permanganate ion - a kinetic study. Inorg. Chem. 1991, 30 (12), 2638-2642.

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(9) Yuan, Z., Van Briesen, J. M. Analysis of biodegradation intermediates ethylenediaminetetraacetate and nitrilotriacetate by high performance liquid chromatography. In Biogeochemistry of Chelating Agents; ACS Symposium Series 910; American Chemical Society: Washington, DC, 2005; pp 139-148. (10) No¨rtemann, B. Biodegradation of chelating agents: EDTA, DTPA, PDTA, NTA and EDDS. In Biogeochemistry of Chelating Agents; ACS Symposium Series 910; American Chemical Society: Washington, DC, 2005; pp 150-170. (11) Kembhavi, M. D. S. R.; Harihar, A. L.; Nandibewoor, S. T. Kinetics of oxidative deamination and decarboxylation of L-asparagine by alkaline permanganate: A mechanistic approach. Inorg. React. Mec. 2001, 3 (1), 39-49. (12) Harihar, A. L.; Kembhavi, M. R.; Nandibewoor, S. T. Kinetics of oxidative degradation of L(+) lysine by alkaline permanganate - A mechanistic approach. Indian J. Chem., A 2000, 39 (7), 769-774. (13) Chougale, R. B.; Panari, R. G.; Nandibewoor, S. T. Kinetics and mechanisms of alkaline permanganate oxidation of L(+)aspartic acid. Oxidat. Commun. 1999, 22 (2), 298-307. (14) Chougale, R. B.; Panari, R. G.; Nandibewoor, S. T. Kinetics and mechanism of alkaline permanganate oxidation of L-glutamic acid. Oxidat. Commun. 1998, 21 (4), 565-573. (15) Jaky, M.; Szeverenyi, Z.; Simandi, L.i. Formation of manganate (V) in oxidations by permanganate ion in strongly alkaline solutions. Inorg. Chim. Acta 1991, 186 (1), 33-37. (16) Duff, M. C.; Hunter, D. B.; Hobbs, D. T.; Jurgensen, A.; Fink, S. D. Characterization of Plutonium, Neptunium, Strontium on Manganese Solids from Permanganate Reduction; Report WSRCTR-2002-00366; Westinghouse Savannah River Company: Aiken, SC, 2002. (17) Jaky, M.; Szammer, J.; Simon-Trompler, E. Kinetics and mechanism of the oxidation of ketones with permanganate ions. J. Chem. Soc. - Perkin Trans. 2 2000, 7, 1597-1602. (18) Panari, R. G.; Harihar, A. L.; Nandibewoor, S. T. Kinetics and mechanism of oxidation of 1,10-phenanthroline by alkaline permanganate. J. Phys. Org. Chem. 1999, 12 (4), 340-346. (19) Pol, P. D.; Mahesh, R. T.; Nandibewoor, S. T. Free radical intervention in the oxidation of nicotinamide by alkaline permanganate - A kinetic study. J. Chem. Res. 2002, 11, 533534. (20) Mulla, R. M.; Nandibewoor, S. T. Mechanistic and spectral investigation of the oxidation of 4-hydroxycoumarin by aqueous alkaline permanganate using the stopped flow technique. Polyhedron 2004, 23 (16), 2507-2513 (21) Sun, Y. F.; Kirschenbaum, L. J.; Kouadio, I. Kinetics and mechanism of the multistep oxidation of ethylenediaminetetraacetate by Ag(OH)4- in alkaline media. J. Chem. Soc., Dalton Trans. 1991, 9, 2311-2315. (22) So¨rensen, M.; Zurell, S.; Frimmel, F. H. Degradation pathway of the photochemical oxidation of ethylenediamine tetraacetate (EDTA) in the UV/H2O2 process. Acta Hydrochim. Hydrobiol. 1998, 26 (2), 109-115. (23) Noradoun, C. E.; Cheng, I. F. Kinetics of EDTA degradation induced by oxygen activation in a zerovalent iron/air/water system. Environ. Sci. Technol. 2005, 39 (18), 7158-7163. (24) Stewart, R. Oxidation by Permanganate. In Oxidation in Organic Chemistry Part A; Academic Press: New York, 1965. (25) Rosenblatt, D. H.; Davis, G. T.; Hull, L. A.; Forberg, G. D. Oxidations of amines. V. Duality of mechanisms in the reaction of aliphatic amines with permanganate. J. Org. Chem. 1968, 33 (4), 1649-1650. (26) Dennis, W. H., Hull, L. A.; Rosenblatt, D. H. Oxidations of amines. IV. Oxidative fragmentation. J. Org. Chem., 1967, 32 (12), 37833787. (27) Hull, L. A.; Davis, G. T.; Rosenblatt, D. H.; Mann, C. K. Oxidations of amines. VII. Chemical and electrochemical correlations. J. Phys. Chem. 1969, 73 (7), 2142-2146. (28) Wiberg, K. B.; Stewart, R. The mechanisms of permanganate oxidations. I. The oxidation of some aromatic aldehydes. J. Am. Chem. Soc. 1955, 77 (7), 1786-1795. (29) Vivekanandan, S.; Venkatarao, K.; Santappa, M.; Shanmuganathan, S. Interactions of alkaline permanganate with chloramine T - a kinetic study. Indian J. Chem., A 1983, 22 (3), 244-245.

Received for review March 7, 2006. Revised manuscript received June 8, 2006. Accepted June 13, 2006. ES0605366