Environ. Sci. Technol. lSS3, 27, 409-412
(12) Jersey, J. A. Ph.D. Dissertation, University of North Carolina, Chapel Hill, NC, 1991. (13) Yoon, J.; Jensen, J. N. Presented a t the AWWA WQTC, Nov 1992. (14) Morris, J. C.; Isaac, R. A. In Water Chlorination: Environmental Impact and Health Effects; Jolley, R. L., et al., Eds.; Ann Arbor Science: Ann Arbor, MI, 1983; Vol. 4, Book 1, pp 49-62. (15) Jensen, J. N.; Johnson, J, D. Environ. Sci. Technol. 1990, 24, 985-990. (16) Box, M. J. Comput. J. 1965, 8 ( l ) , 42-52. (17) White, G. C. The Handbook of Chlorination,2nd ed.; Van Nostrand Reinhold Co.: New York, 1986. (18) Isaac, R. A.; Morris, J. C. Environ. Sci. Technol. 1985,19, 810-814. (19) Morris, J. C. In Principles and Applications of Water Chemistry;Faust, S. D., Hunter, J. V., Eds.; John Wiley and Sons, Inc.: New York, 1967; pp 23-53. (20) Carey, F. A.; Sundberg, R. J. Advanced Organic Chemistry. Part A., 3rd ed., Plenum Press: New York, 1990; p 286. (21) Gray, E. T., Jr.; Margerum, D. W.; Huffman, R. P. In
were found to be analogous. Both processes were affected by the chloramine (or parent amine) basicity to a similar degree.
Literature Cited (1) Wolfe, R. L.; Ward, N. R.; Olson, B. H. Environ. Sci. Technol. 1985, 19, 1192-1195. (2) Feng, T. H. J.-Water Pollut. Control Fed. 1966, 16, 614-628. (3) Weil, I.; Morris, J. C. J. Am. Chem. SOC. 1949,71,1664-1671. (4) Friend, A. G. Ph.D. Dissertation, Harvard University, Cambridge, MA, 1956. (5) Margerum, D. W.; Gray, E. T., Jr.; Huffman, R. P. In
Organometals and Organometalloids;Occurrence and Fate in the Environment;Brinckman, F. E., Bellama, J. M., Eds.; ACS Symposium Series 82; American Chemical Society: Washington, DC, 1978; pp 278-291. (6) Granstrom, M. L. Ph.D. Dissertation, Harvard University, Cambridge, MA, 1954. (7) Hussain, A,; Higuchi, T.; Hurwitz, A.; Pitman, I. H. J. Pharm. Sci. 1972,19, 371-374. (8) Snyder, M. P.; Margerum, D. W. Inorg. Chem. 1982,21, 2545-2550. (9) Isaac, R. A,; Morris, J. C. Environ. Sci. Technol. 1983,17, 738-742. (10) Lukasewycz, M. T.; Bieringer, C. M.; Liukkonen, R. J.; Fitzsimmons, M. E.; Corcoran, H. F.; Lin, S.; Carlson, R. M. Environ. Sci. Technol. 1989, 23, 196-199. (11) Standard Methods for the Examination of Water and Wastewater,17th ed.; American Public Health Association, American Water Works Association, Water Pollution Control Federation: Washington, DC, 1989.
Organometals and Organometalloids: Occurrence and Fate in the Environment; Brinckman, F. E., Bellama, J. M., Eds.; ACS Symposium Series 82; American Chemical Society: Washington, DC, 1978; pp 264-277. (22) Isaac, R. A.; Morris, J. C. In Water Chlorination: Environmental Impact and Health Effects; Jolley, R. L., et al., Eds.; Ann Arbor Science: Ann Arbor, MI, 1983;Vol. 4, Book 1, pp 63-75. Received for review June 26, 1992. Revised manuscript received October 6, 1992. Accepted October 19, 1992.
COMMUNICATIONS Formation of Carbon Monoxide from the Photodegradation of Terrestrial Dissolved Organic Carbon in Natural Waters Richard L. Valentine"
Department of Civil and Environmental Engineering, University of Iowa, Iowa City, Iowa 52242 Richard 0. Zepp
Environmental Research Laboratory, U.S. Environmental Protection Agency, Athens, Georgia 30613
Introduction Photochemical degradation of dissolved natural organic matter results in the formation of a number of products (1-4) and is a process which may govern the rate of carbon turnover in the ocean (3, 4). One major volatile photoproduct of the dissolved organic matter (DOM) in the sea is carbon monoxide, a gas that is emitted to the atmosphere (4-6). Although not a greenhouse gas, carbon monoxide has an important effect on the radiative balance of the atmosphere through its enhancement of the buildup of methane, ozone, and other radiatively important atmospheric trace gases (7, 8). This effect results mainly through reactions of carbon monoxide with the hydroxyl radical, a key intermediate in the complex set of chemical reactions occurring in the troposphere (7,8). Great uncertainty exists regarding the source strength of CO derived from the photodegradation of DOM. Past 0013-930X/93/0927-0409$04.00/0
studies have focused on open ocean water (3-6). A recent study (6) of CO fluxes in the Pacific Ocean suggests that the oceanic CO flux may be as high as 200 Tg/y. Modeling and measurements indicate that this marine source is a significant fraction of the CO emissions in the Southern Hemisphere (6, 9). Wetlands and near-coastal regions are also potentially important sources of CO. Wetland waters have considerably greater near-surface photoreactivity than other natural waters due to their higher DOC content and much stronger absorption of sunlight ( I , 10-12). Biologically refractory terrestrial DOC also has high photochemical reactivity (1,10-12). This suggests that production of CO in wetlands may be important to both the global and the regional atmospheric budgets of CO given that 5.3 million km2 of wetlands exists (13). In addition, terrestrial DOM is a major source of organic matter in near-coastal regions
0 1993 American Chemical Society
Environ. Sci. Technol., Vol. 27, No. 2, 1993 409
Table I. Samples Studied and Photoproduction Rates of Carbon Monoxide from Exposure to Simulated Solar Radiation approximately Equal to the Annual Average Value at 40' N Latitude
source
location
Kinoshe Lake Houghton Marsh Okefenokee Swamp Suwannee River Intracoastal Waterway Live Oak Oyster River fulvic Houghton Marsh" Houghton Marshb Suwannee River" Suwannee River' soil fulvic Contech fulvic Fluka humic
51' N, 81' W 41' N 87" W 31" N, 83" W 29' N, 83' W 30' N, 81.5O W 30' N, 83.5' W 43' N, 71° W 41" N, 81" W 41" N, 87' W 29" N, 83" W 29' N, 83' W 43' N, 11' W
DOC (mg of C/L) 15 31 31 41 3.2 7.1 56 28 24 33 21
56 52 30
near-surface CO prodn rate divided by a35$ (nmol m L-l h-l)
est annual av daily near-surface C turnover rate at 40" N ( % C/d)
260 1330 1200 1500 110
12
0.2 0.3
90
11
3200 480 120 490 41 3330 1600 1500
18 14
absn coeff near-surface CO ( a 3 d dat prodn rate 350 nm (m-l) (nmol L-' h-l) 21 88 92 120 6.2 1.6 170 34.5 10.4 36.8 3.1 190 230 180
14 14 13 17
0.3
12
0.3 0.3 0.1 0.5 0.2 0.05
13
0. I
11
0.01 0.5
77 7
0.3
8
0.5
"Faded by irradiating 50 h prior to CO production rate determination. bFaded by irradiating 119 h prior to CO production rate determination. 'Faded by irradiating 194 h prior to CO production rate determination. dValues of u350= 2.303A3,,/b, where A350is absorbance and b is Dath length in meters.
(14, 1.9, which comprise another 49 million km2 of the Earth's surface (16). In this study we investigated the photochemical formation of 60 in water samples obtained from wetlands, lakes, and near-coastal/shelf areas and in aqueous solutions of soil organic matter. All of these samples contain DOM that is largely derived from terrestrial sources. Our studies show that, although the water samples had widely varying optical properties and CO photoproduction rates, the efficiencies for photochemical CO formation were remarkably similar in all waters examined. Model calculations further indicate that photodegradation of terrestrial DOM (e.g., in wetland and near-coastal environments) may be an important global source of carbon monoxide and a key process in the cycling of DOM in these environments.
Materials and Methods Natural water samples were obtained that were representative of wetlands and high-DOC lakes from the boreal forest/taiga regions of North America (Kinoshe Lake, Ontario, Canada; Houghton Marsh, Michigan) and a southern wetland ecosystem (Okefenokee Swamp and Suwannee River, Georgia). Two coastal water samples also were studied, one from the Intracoastal Waterway off the Atlantic coast of Florida, and the other obtained near Live Oak Island off the Gulf Coast of Florida. Studies also used prepared solutions of fulvic acids extracted from a soil and river in New Hampshire that were described in a previous study (17) and two commercially available materials, Contech fulvic acid and Fluka "humic acid", which were dissolved in deionized water without pH adjustment and stirred for 48 h. The DOC content and the absorption coefficients at 350 nm (a35o),a wavelength where light absorption is dominated by humic substances in water (4, 18),varied over 1order of magnitude in the waters (Table I). All water samples were air saturated and filtered though 0.2-pm filters to remove particulates prior to irradiation. The pH of all water samples was between approximately 5.5 and 6.5. Additional studies showed that CO production was not a function pH over this range. Comparisons indicated that filtration had no detectable effect on the carbon monoxide photoproduction rates in the samples. Quantum yields (ratio of moles of CO produced to einsteins of light absorbed) were determined using a Kra410
Environ. Sci. Technoi., Vol. 27, No. 2, 1993
tos/Schoeffel 1 KW Reaction Chemistry System with monochromator set at 5.0-mm bandwidth. Samples were held in gas-tight quartz cells with 1-13-cm path lengths. Light intensity was measured using ferrioxalate (19) and p-nitroanisole (20) actinometers. CO formation rates were determined for samples exposed to simulated solar radiation from a Spectral Energy Model LH153 solar simulator (65 mW/cm2). These samples were diluted with deionized water so that the absorbance at 350 nm was less than 0.05 cm-l to eliminate artifacts caused by light attenuation and contained in 13.5 cm X 1.5 cm round quartz tubes. Reported "near-surface" CO formation rates were determined by multiplying observed rates by the dilution factor. Irradiance (200-3000 nm) was measured using an International Light Model IL 1700 radiometer in these solar simulator experiments. In addition, p-nitroanisole actinometers (20) were used to verify that the UV light intensity was constant over the course of the irradiations. No significant formation or consumption of CO was observed in dark controls. Even at 50 "C, thermal CO formation was only 1% of the photochemical production rate. Studies examining the relationship between CO formation rate and extent of fading were conducted using Houghton and Suwannee River samples. Fading was accomplished by exposing undiluted samples contained in the quartz tubes to continuous simulated solar radiation over a 194-h time period. Sample tubes were periodically opened and samples withdrawn for absorbance measurements and determination of near-surface CO formation rates as previously described. Carbon monoxide was determined in headspace gas samples using a Perkin-Elmer Model 2000 gas chromatograph fitted with a methanizing flame ionization detector. CO measurements were calibrated versus secondary standard gas mixtures obtained from Scotty Specialty Gases. The Scotty standards were calibrated versus CO standard gas mixtures obtained from the National Institute of Standards and Technology. Total CO produced was determined through use of Henry's constant to relate the headspace concentration to water concentration. Absorbance measurements were made using a Shimadzu Model 265 scanning spectrophotometer. Analysis of p-nitroanisole was performed using a Waters Associates Model 6000 HPLC and OD$-2 column with 60:40% acetonitrilewater as the mobile phase and a Applied Biophysics UV-visible detector. Fe(I1) was analyzed using a modified
2.0
10" A
-
1.6
k l "E
m 1.2 -
v
3 X
U
200
'
I
300
'
I
400
'
I
'
I
500 600 Wavelength (nm)
'
700
Flgure 1. Quantum yields for carbon monoxide formation as a functlon of wavelength for several natural water samples. Samples were filtered through 0.2-pm filter and alr saturated.
ferrozine method (21). DOC was determined using a Dohrman Model DC-60 carbon analyzer. Results and Discussion The near-surface production rates (Table I), defined as the rates obtained in optically thin systems (absorbance
