Environ. Sci. Technol. 1987, 21 443-450 I
(10) Millero, F. J. Limnol. Oceanogr. 1986, 31(4), 839-847. (11) Avrahami, M.; Gelding, R. M. J . Chem. Soc. A 1968, 647-651. (12) Chen, K. Y.; Morris, J. C. Environ. Sei. Technol. 1972,6, 529-537. (13) Hoffman, M. R.; Lim, B. C. Enoiron. Sei. Technol. 1979, 13, 1406-1414. (14) Hoffman, M. R. Enuiron. Sei. Technol. 1977, 11, 61-66. (15) O'Brien, D. J.; Birkner, F. G. Environ. Sei. Technol. 1977, 11, 1114-.1120. (16) Sorokin, Y. L. Okeanologiya (Moscow)1971,11,423-431. (17) Skopintsev, B. A.; Karpov, A. V.; Vershinina, 0. A. Sou. Oceanogr. 1964,4, 55-73. (18) Cline, J. D. PbD. Thesis, University of Washington, Seattle, WA, 1969. (19) Ostlund, G. H.; Alexander, J. J. Geophys. Res. 1963, 68, 3995-3997. (20) Almgren, T.; Hagstrom, L. Water Res. 1974,8, 395. (21) Garnett, S.; Millero, F. J., unpublished results, 1984. An
products formed as a function of temperature and ionic strength. These measurements will hopefully lead to a better understanding of the mechanism of the oxidation. Registry No. H2S, 7783-06-4.
Literature Cited Emerson, S.; Jacobs, L.; Tebo, B. in Trace Metals in Seawater;Wong, C. S.; Burton, J. D.; Bruland, K.; Goldberg, E., Eds.; Plenum: New York, 1983; pp 579-608. Boulegue, J.; Lord, C. J., 111; Church, T. M. Geochim. Cosmochim. Acta 1982,46, 453-464. (3) Millero, F. J. Mar. Chem. 1986, 18, 121-147. (4) Cline, J. D.; Richards, F. A. Environ. Sei. Technol. 1969, 3 838-843.
Cline, J. D. Limnol. Oceanogr. 1969, 14, 454-458. Benson, B. B.; Krause, D., Jr.; Peterson, M. A. J . Solution Chem. 1979,8,655-690. Benson, B. B.; Krause, D., Jr. Limnol. Oceanogr. 1984,29,
ORION sulfide electrode was used in these studies.
620-632.
Harned, H. S.; Owen, B. B. in The Physical Chemistry of Electrolytic Solutions; Reinhold: New York, 1958; p 736. (9) Covington, A. K.; Bates, R. G.; Durst, R. A. Pure Appl. Chem. 1985,57, 531-542.
Received for review July 9, 1986. Accepted December 1, 1986. This study was supported by the Oceanographic Section of the National Science Foundation and the Officeof Naval Research.
Nitrate-Induced Photooxidation of Trace Organic Chemicals in Water Richard G. Zepp" Environmental Research Laboratory, US. Environmental Protection Agency, Athens, Georgia 30613
Jurg Holgnd" and Helnz Bader Swiss Federal Institute for Water Resources and Water Pollution Control (EAWAG), CH-8600 Dubendorf, Switzerland
The oxidation kinetics of butyl chloride, nitrobenzene, anisole, and methylmercury in the presence of hydroxyl radical ('OH) scavengers were used to determine the rate and quantum efficiency for production of 'OH from irradiated nitrate ions in water. The experiments were conducted under steady-state irradiations with monochromatic radiation (313 nm) and with sunlight. The mean quantum efficiency for 'OH production at 313 nm rises from 0.013 f 0.002 at 20 OC to 0.017 f 0.003 at 30 "C in the pH range 6.2-8.2. Results of this study are used to estimate nitrate-induced photooxidation rates of trace organic chemicals under a variety of environmental conditions.
t llght
-
NO3-*
followed by
'0-
+ H,O
-
TT
N02NO2
+ O(3P) + '0-
*OH + OH-
(1)
(2)
(3)
'0- is rapidly protonated to its conjugate acid, the hydroxyl radical (eq 3), a potent oxidant that reacts much more rapidly with most organic chemicals than does atomic oxygen, O(3P) (6). The major fate of the atomic oxygen produced in reaction 1 is likely to be reaction with dioxygen to form ozone. The ozone is rapidly consumed by reaction with NOzor by decomposition to 'OH (8). The possible role of nitrate photolysis in producing hydroxyl radicals in seawater was first recognized by Zafiriou (9). On the basis of analysis of earlier studies by Hamilton (IO), Zafiriou concluded that nitrate photolysis is likely to provide a significant source of free radicals in the sea. Subsequent studies by Zafiriou and True ( I I ) , however, indicated that nitrate photolysis was much slower than reported by Hamilton. The reason(s) for the difference in these results is (are) unclear. Reoxidation of the reaction products to nitrate under the conditions of Zafiriou's experiments seems to be the most plausible explanation. Russi et al. (2) demonstrated through product studies that photolysis of nitrate in freshwater samples produced hydroxyl radicals. The kinetic results in this study were not designed in such a way that they could be generalized to other natural waters. Nonetheless, the results did show that nitrate-induced photooxidations are significantly rapid in certain natural waters. Haag and Hoign6 (12) suggested
(a,
Introduction
The nitrate in aquatic environments has long been known to be involved as an electron acceptor in the biological oxidation of organic substrates (I). Recent evidence indicates that nitrate ions also promote the photochemical oxidation of trace organic compounds in water. Russi and co-workers (2) have provided kinetic studies of the nitrate-induced photooxidation of benzene in distilled and natural water samples. Other studies have demonstrated that nitrate ions enhance photooxidation rates of methylsilyl derivatives (3) as well as other commercially important organic chemicals ( 4 ) . The basic photochemistry of nitrate ions in water has been reviewed by Wagner and co-workers (5). The irradiation of nitrate in its long-wavelength absorption band (maximum 302 nm) results in two primary photochemical processes:
* Author to whom correspondence should be addressed. 0013-936X/87/0921-0443$01.50/0
N03-
0 1987 American Chemical Society
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from their studies on the photolytic production of OH radicals in a lake water that nitrate could well be the dominant precursor. Here we report kinetic studies involving steady-state irradiations that are designed to complement and clarify the past studies. The main objective was to determine generalizable results that can be used to reliably estimate the effects of nitrate ions on organic photooxidation rates in natural waters. Oxidation kinetics of selected trace organic chemicals were analyzed to determine the quantum yield for OH radical production from nitrate photolysis with light of wavelengths characteristic of sunlight. The results then were used to estimate typical rates of nitrate-photoinitiated OH radical oxidations as a function of natural water composition. This approach permits usage of the extensive compilations (13) of kinetic data on reactions of OH radicals with organic chemicals in water. The concepts used here are discussed in more detail by Hoign6 et al. (8, 14, 151, Haag and Hoign6 (12) (radical kinetics), and Zepp and Cline (16) (photochemical kinetics).
Experimental Section Materials. Reagent-grade butyl chloride, sodium nitrate, and l-octanol were obtained from Merck AG. Nitrobenzene (Chem Service) and anisole from Fluka were of analytical grade. These chemicals were used as received. p-Nitroanisole was obtained from Merck AG and purified by recrystallization from 95% ethanol. Methylmercury was prepared as described earlier (14). The water used was first deionized and then doubly distilled, with a third distillation from alkaline permanganate under a nitrogen atmosphere. Equipment. Irradiations were performed in merrygo-round reactors at both EAWAG in Switzerland and at the Environmental Research Laboratory, Athens, in the U.S. (12, 17). Both reactors used filtered light from high-pressure mercury lamps as the light source. The 313 line was isolated with an aqueous solution of 1 mM potassium chromate in 2.3% potassium carbonate and a borosilicate glass sleeve as the filter. In some experiments, only the borosilicate glass sleeve was used as filter. The temperature was maintained at 20.0 f 0.5 or 30.0 f 0.1 "C in the reactor; the former was used in most experiments. Experiments also were conducted with a Schoeffel reaction chemistry system to provide monochromatic radiation (17). Irradiations in sunlight were conducted in quartz tubes arranged in a rack. Analyses were performed with a Waters Associates Model 6000 high-pressure liquid chromatograph (HPLC) and a Micromeritics 7170 HPLC; both were equipped with variable-wavelength UV detectors. Analyses of the butyl chloride solutions employed a gas chromatograph equipped with a FID detector. A 2 m X 4 mm (i.d.) glass column packed with 2% SE-30 on acidwashed DMCS-treated Chromosorb G (SO/lOO mesh) was operated isothermally at 100 OC. Electronic absorption spectra were obtained on a Uvikon Model 810 UV spectrophotometer. The mineralized methylmercury was determined by the Hatch/AAS method as described elsewhere (14). General Procedures. Air-saturated stock solutions containing phosphate buffer (0.005 M) at pH 6.2,8.2, or 8.0,l-octanol (ranging from 15 to 300 pM), nitrate (0.2-4.0 mM), and reactant (0.5-4 pM) were prepared in distilled water. Actinometer solutions consisted of p-nitroanisole (10 pM) in distilled water (18). Aliquots of the reactantlnitrate solution and the actinometer solutions were simultaneously irradiated in the merry-go-round or in sunlight (12, 17). 444
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Except in the case of methylmercury, internal standards were added following irradiation. These solutions were analyzed by HPLC (anisole, nitrobenzene, p-nitroanisole) or by headspace gas chromatography (GC). p-Nitrotoluene was used as internal standard for the liquid chromatographic analyses. In all cases, the HPLC analyses were performed on an ODs-2 column with 60%-40% acetonitrile-water as the mobile phase. Three wavelengths were used in the UV detector: 220 nm, anisole; 267 nm, nitrobenzene; 315 nm, nitroanisole. Butyl chloride analyses were conducted with the headspace GC method (12). The procedure used to compute light intensity from the nitroanisole actinometer kinetic data was discussed by Zepp et al. (17, 19). Temperature studies showed that the quantum yield (313 nm) for this actinometer ranged from 3.2 X lo4 at 20 "C to 2.5 X lo4 at 30 OC. Nitrite formation was determined by the sulfanilic acid method on a Technicon analyzer. Photolysis rate constants of nitrate were calculated with a modified version of a computer program that is described by Zepp and Cline (16). The current version of this program, GCSOLAR, uses equations developed by Green, Cross, and Smith (20) to compute the solar UV irradiance. This program, which is available in PC-compatible form, is available on request from R.G.Z.
Results and Discussion Kinetic Model. Previous studies discussed above (eq 1-3) indicate that the oxidation of organic compounds is photoinitiated at wavelengths >290 nm by nitrate ions through the following reactions involving the intermediacy of the hydroxyl radical: ,OH
Products products
where M and Si represent solutes that rapidly react with 'OH with rate constants kM and hi,respectively. In this paper M is used to symbolize some readily detectable trace substance that was used as a probe for hydroxyl radicals, and Si represents other more concentrated solute(s) that scavenge 'OH in competition with M. If the concentration of M, [MI, is sufficiently low that k ~ [ M is l
