Environ. Sci. Technol, 1983, 17, 107-112
nutrients, either inorganic or organic. Mineralization of low concentrations of organic chemicals may be a result of activities of oligotrophic populations. The finding that propionate or glucose at 38.5 pg of carbon/mL markedly inhibited the mineralization of phenol is consistent with the view that these organisms are sensitive to organic chemicals at concentrations often used to test for biodegradation; that is, at levels that would sustain eutrophic microorganisms. Alternatively, the cells may have selectively utilized glucose and propionate rather than the small amount of'phenol. It has been shown previously that mineralization of 2,4-D in fresh waters sometimes ceased at concentrations as low as 200 ng of 2,4-D/mL, presumably because of the sensitivity of the active organisms (8). The significance of the oligotrophs, which have been reported in natural waters (9-11), thus requires further inquiry. Concentrations of substrates below which no mineralization occurs have been demonstrated in some natural waters (6)and in pure culture (12). The failure to observe such thresholds in other waters (8) may result from the presence in these waters of nutrients that influence the indigenous community. In support of this view is the finding of Law and Button (13) that the threshold concentration for glucose in culture was lowered in the presence of arginine and was lowered beyond the limits of detection in the presence of a mixture of amino acids. The especially marked enhancement by arginine of the mineralization of less than 2 pg of phenol/mL by the community from lake water may reflect a similar type of microbial response. A threshold may be expected in environments in which the indigenous populations are obligate eutrophs that are not able to metabolize the test chemicals at low levels regardless of the presence of other nutrients.
Acknowledgments We thank Deidre M. Brophy for technical assistance. Registry No. p-Nitrophenol, 100-02-7;glucose, 50-99-7; adenine, 73-24-5; arginine, 74-79-3; propanoic acid, 79-09-4; Mg, 7439-95-4; NH4, 14798-03-9; Ca, 7440-70-2; Fe, 7439-89-6; Mo, 7439-98-7; Mn, 7439-96-5; phenol, 108-95-2.
Literature Cited (1) Tanaka, N.; Ueda, Y.; Onizawa, M.; Kadota, H. Jpn. J. Limnol. 1977,38, 41. (2) Vaccaro, R. F. Limnol. Oceanogr. 1969, 14, 726. (3) Nesbitt, H. J.; Watson, J. R. Water Res. 1980, 14, 1689. (4) Alexander, M. In "Microbiology 1980"; Schlesinger, D., Ed.; American Society for Microbiology: Washington, DC, 1980; p 328. (5) Nesbitt, H. J.; Watson, J. R. Water Res. 1980, 14, 1683. (6) Boethling, R. S.; Alexander, M. Appl. Environ. Microbiol. 1979,37, 1211. (7) Subba-Rao, R. V.; Rubin, H. E.; Alexander, M. Appl. Environ. Microbiol. 1982, 43, 1139. (8) Rubin, H. E.; Subba-Rao, R. V.; Alexander, M. Appl. Environ. Microbiol. 1982, 43, 1133. (9) Akagi, Y.; Taga, N.; Simidu, U. Can. J . Microbiol. 1977, 23, 981. (10) Kuznetsov, S. I.; Dubinina, G. A.; Lapteva, N. A. Annu. Rev. Microbiol. 1979, 33, 377. (11) Moaledj, K.; Overbeck, J. Arch. Hydrobiol. 1980,89, 303. (12) Jannasch, H. W. Limnol. Oceanogr. 1967,12, 264. (13) Law, A. T.; Button, D. K. J. Bacteriol. 1977, 129, 115.
Received for review December 28, 1981. Revised manuscript received August 9, 1982. Accepted October 26, 1982. This research was supported by the Environmental Protection Agency cooperative agreement CR806887. The statements do not necessarily reflect the views and policies of the Environmental Protection Agency.
Environmental Fate and Effects of Ethylene Oxide Richard A. Conway," Gene 1.Waggy, Milton H. Splegel, and Ronald L. Bergiund Research and Development Department, Solvents and Coating Materials Division, Union Carbide Corporation, South Charleston, West Virginia 25303
rn A study has been completed to learn more about the fate and effects of ethylene oxide (EO) in the environment and especially to fill gaps revealed in a recently published in-depth survey of available literature. Key findings on this large-volume product (3 X loe kg/year) include (1) desorption coefficient of 0.36 times that of oxygen, (2) hydrolysis half-life of 14 days in natural fresh water at 25 "C, (3) hydrolysis/hydrochlorination half-life of 9 days in 3% salt water at 25 "C and a chlorohydrin/glycol product ratio one-tenth of literature estimate, (4) biooxidation of EO and derivatives at a rate indicating nonpersistence, and (5) aquatic toxicity at moderate levels (e.g., 96-h LCmwith fathead minnows of 84 mg/L, 48-h LC50 with Daphnia magna of 212 mg/L, and 48-h LC50 with brine shrimp of 745 mg/L). Also, literature cited indicates that in air ethylene oxide is nonpersistent due to washout by rain and degradation by free-radical processes.
Introduction The production of ethylene oxide (EO) in the United States is about 3 X lo9 kg/year, ranking it among the top 15 chemicals in volume. After various governmental committees identified a need to know more about the environmental fate and effects of this material (1-3), a com0013-936X/83/0917-0107$01.50/0
prehensive paper study toward this end was conducted by Syracuse Research Corp. (SRC) ( 4 ) . Union Carbide Corp. now has completed environmental chemistry and ecotoxicity studies to supplement and extend the SRC findings. Studies were conducted in the areas of volatilization, hydrolysis, hydrochlorination, biodegradation, and aquatic toxicity, As a result, risk analyses can be made for various ecological situations.
Methods EO has a molecular weight of 44.053, freezing point of -112.44 "C, boiling point of 10.5 "C at 760 torr, vapor pressure of 1305 torr at 25 "C, and complete solubility in water (4). The high solubility indicates a low potential to accumulate in fatty tissue or to partition onto clays. The relative desorption coefficient for use in estimating volatilization rate was determined experimentally for EO by using an unbaffled 2-L vessel, 10 cm deep. EO was added to water (22 "C) and rapidly mixed for a 4-h period by using a 3-cm, 45" propellor at 450 rpm. Prior to introducing EO, the oxygen-transferrate was determined for this system. Tests were conducted both with a 5 m/s wind flow over the liquid surface and with no induced wind flow. After addition of EO, the dissolved EO concentration was
0 1983 Amerlcan Chemical Society
Environ. Sci. Technol., Vol. 17,No. 2, 1983 107
determined a t various times by using direct aqueous injection into a gas chromatograph, Hydrolysis of EO was examined in distilled water, a natural fresh water, and salt water. A composite of Kanawha River water collected at three points in Charleston and South Charleston, WV, on Oct 15,1979, was used as the natural fresh water. For the salt water, first a pH 7 buffer was prepared by mixing 30 mL of 0.06 M NaH2P04 with 61 mL of 0.067 M K2HP04,diluting it 10-fold, and sterilizing it by filtration through a 0.22-pm filter. Buffered solutions containing about 200 mg/L EO and 0% , 1% ,and 3% NaCl were tested at 25 "C. EO was added as follows. A cylinder containing EO gas was placed in a hood and connected to a swagebck tee equipped with a septum. Gas flowed through the tee to a rotameter and was vented to the hood. A precalculated volume of ethylene oxide, depending on the temperature and pressure, was collected in a syringe and injected into a septum-sealed, sterile vial containing a known quantity of sterile water buffered at the desired pH. Each vial was almost completely filled to lessen volatilization. Various concentrations of ethylene oxide were prepared to establish an analytical base. The vials were placed in a temperature-controlled, shaken water bath, covered to exclude light. Samples were removed at approximately regular intervals during a 350-h test with a 10-pL syringe. The needle was flamed prior to insertion in the vial if the vial were to be used for more than one sample. The samples were then analyzed to determine the EO concentration. The change in EO reactant concentration with time was determined by gas chromatography with use of 6.1 m X 3.2 mm stainless steel tubing packed with a 30% solution of Tergitol nonionic surfactant TMN-3% sodium methylate on Chromosorb W, NAW, 60/80 mesh. Maximum upper temperature for conditioning is 75 "C. Operating conditions were as follows: injection port, 60 "C; column, 60 "C; a flame ionization detector, 200 "C; helium flow, 22 cm3/min. A sample size of 5 pL was used in a solvent flush technique. The approximate retention time was 6.3 min. A different gas chromatographic column was used for salt water tests, which took place several months later. The column was recommended as possibly satisfactory for both EO and chlorohydrin with regard to accuracy and reasonable elution times. The assumption was made that all noted EO losses were due to the production of ethylene glycol and chlorohydrin; therefore, ethylene glycol could be calculated by difference. The column used was 6.1 m x 3.2 mm stainless steel tubing packed with super PAK20M. Operating conditions were the same as those used for fresh water. Theoretical oxygen demand was calculated on the basis of the oxygen required to completely oxidize the chemical to its lowest energy state, Le., C 0 2 and H2O. The value may be determined by the chemical oxygen demand (COD) wet chemical procedure (chromic acid depletion) published in ref 5. Standard potassium hydrogen phthalate (0.4251 g/L) solutions were measured with each COD series. Nineteen determinations on the standard (theoretical concentration of 500 mg/L) gave a mean of 504 mg/L and standard deviation of 6 mg/L. Biochemical oxygen demand (BOD) was measured by using procedures that also generally follow ref 5. Method changes involved extending the test period to 20 days; maintaining three seed control bottles (2 mL of domestic sewage/ bottle) through the test period for calculating seed oxygen demand; accomplishing reaeration, if needed, by 108
Environ. Sci. Technol., Vol. 17, No. 2, 1983
dividing the BOD bottle contents between two BOD bottles, sealing, shaking 20 times and returning to the original BOD bottle, reading the oxygen level, resealing, and returning to an incubator (6). Standard solutions of glucose/monosodium glutamate (theoretical of 215 mg/L) were determined with each test series. Eighteen measurements gave a mean of 215 mg/L with standard deviation of 3 mg/L. Fish and Daphnia toxicity data were collected by using procedures as published by EPA and ASTM (7). Definitive tests used 10 fish/test concentation in test volumes of 10-15 L. Range-fmding procedures used 2 minnows/test concentration. Only minimal aeration or oxygen blankets were used to maintain sufficient dissolved oxygen in most test series. The oxygen-blanket technique involves raising the dissolved oxygen concentration in the test solution to 15-20 mg/L with a quick sparging with pure oxygen and sealing the contents under a highly oxygenated atmosphere. Dechlorinated (carbon-treated) Charleston tap water was used as dilution water in all tests. Brine shrimp test procedures were described by Price et al. (6). For bacterial toxicity tests, selected concentrations of the test material were incubated for 16 h at 22 "C on a shaker in the presence of nutrients, buffers, growth substrate, and sewage microorganisms. Toxicity is indicated when the resulting turbidity is less than 50% of the control (8).
Results and Discussion Kinetics of Desorption from Water Bodies to the Atmosphere. The transfer of an organic across a gasliquid interface will follow first-order kinetics and can be represented by the equation dC/dt = -Kd(C - X/H) (1) where C and X are the concentrations of the organic in liquid and gas phases, respectively, H is Henry's constant (atm/mol fraction), and Kd is the desorption coefficient. Henry's constant can be calculated from the activity coefficient, y ((mol fraction)-'), and the vapor pressure, P, (atm), from the relationship H = yP,. In most cases other than diffused aeration systems C >> X / H therefore, this equation can be conveniently expressed in the form dC/dt = -Kdc (2) According to the classical two-film mass transfer model, the desorption coefficient can be presented in terms of the gas and liquid film mass transfer coefficients: 1/Kd = 1 / K l + l/(HKg) (3) where Kl and Kgare the liquid and gas film mass transfer coefficients, respectively. With utilization of the concept proposed by Liss and Slater ( 9 ) ,initially applied by Mackay and Leionen (10) and others (11, 12), the desorption coefficient of one chemical in a system can be estimated from the known value of a "tracer" chemical such as oxygen: &(X)
= Kd(0,)[D(X)/D(0,)I,o'6/(1 K,/HKg) (4)
where D(x) and D(0,) are the diffusion coefficients in water for the organic and oxygen, respectively. The exponent 0.5 for the diffusion coefficients, derived from the penetration theory, is consistent with data reported by Sweeris (13) and Smith and Bomberger (14). Smith et al. (15)in a recent computer reduction of their data for high volatility chemicals found that an exponent of 0.61 (film-penetration theory) better fit their results.
I
Table I. Effect of Temperature on Calculated Relative Desorption Coefficient of Ethylene Oxide
mol haction-Ib
H, atm (mol fraction)
5.4 5.5 -5.5
5.3 7.9 11.2
7,
"C
&/K,
P,,a atm
10 20 30
9.1 11.5 14.5
0.95 1.44 2.03
T,
a
Reference 17.
Reference 18.
I
1
1
1
DE01 DO:
ad
0.7 0.7 0.7
0.31 0.34 0.36
-I ETHYLENE OXIDE REMA I N 1NG 5 M/SEC WIND
Reference 19.
Table 11. Desorption of Ethylene Oxide from Watera 1
450 450 a
0 5
2.0 2.26
0.72 0.88
0.36 0.39
MIXING TIME,
HOLRS
Flgure 1. Desorption rate of ethylene oxide from water.
Calculated a d ( 2 2 "C) = 0.34.
However, in their analysis, they included a number of organics not completely controlled by the liquid film resistance (IO,12). Because of the difficulty of measuring experimentally the gas-film mass transfer coefficient (26), estimates have generally been limited to volatile organics (H >> Kl/Kg). However, the results of Liss and Slater (9) and Dilling (12) suggest that the ratio Kl/Kg might reasonably be represented as a constant (temperature dependent), equal to 11.5 atm/mol fraction. Thus, for any chemical, a relative desorption coefficient, ad, can be defined according to the relationship
STERILE D I S T I L L E D WATER
t
LOG RIVER WATER STERILE RIVER WA,lER STERILE D I S T I L L E D WATER
la1
K lSEC-ll
tlALF-LIFE. DAYS
-6.25 -6.21
-6.18
12 .?
where 8 is a temperature correction factor, 1.024, derived 2o 15 I I I I I I J from the relationship K,/K,, proportional to ( T / M ) ~ . ~ , 0 50 100 150 200 250 303 350 where p is the liquid viscosity term. For 10,20 and 30 OC, the calculated a d values for ethylene oxide will be 0.31, T I M , HOURS 0.34, and 0.36, respectively (Table I). Flgure 2. Ethylene oxide hydrolysis in fresh water at pH 7.4 and 25 The relative desorption coefficients experimentally deOC. termined for EO in this system were 0.36 (no wind) and 0.39 (wind) (Figure 1 and Table 11). These values are pH of the natural river water. The hydrolysis results reasonably consistent with a calculated a d of 0.34 at 22 O C indicated normal first-order kinetics in all cases (Figure by using eq 5 (H -8.2 atm/g mol fraction)). The higher 2). There were slight differences in the rate constants for magnitude of the experimental values over the calculated the sterilized samples (half-life of 12.2 and 12.9 days) value could indicate that the Kl/Kgratio decreases with compared to the as is Kanawha River sample (half-life of increasing turbulence (0.36) and wind flow (0.39). How14.2 days), but these were well within the error limits ever, the calculation of a d does give a reasonable assessdiscussed by Mabey and Mill (21). ment of the desorption potential of ethylene oxide from A critical review of published studies comparing the water. These values indicate that ethylene oxide will be rates of hydrolysis of chemicals in natural and pure waters desorbed from a water body with a rate dependent upon shows that the rate constants are the same and that labthe actual oxygen-transfer rate in the system. The rate oratory measurements can provide data for estimating will, however, be less than for volatile organics of low hydrolysis rates in the environment (22).Figure 3, a plot solubility such as benzene, toluene, and chloroform, for of rate constant vs. pH for two temperature levels, shows which Rathburn and Tai (20) reported experimentally the data in perspective and indicates that there were no determined a d values of around 0.65. significant differences. It also shows that water temperAvailable information on epoxide properties collected ature will probably have a greater effect on half-life than by SRC suggests that atmospheric degradation by freeexpected pH differences in natural waters. radical oxidation is relatively fast, which limits the geoKinetics of Salt Water Hydrolysis. The hydrolysis graphic range and time of human exposure hazard (4). of ethylene oxide at pH 7 and 25 "C in a salt-free solution Also washout by rain will reduce levels in air. indicated a first-order reaction with a log k of -6.21 and Kinetics of Fresh Water Hydrolysis. Three samples, a half-life of 314 h (Table 111). These data are in close sterilized distilled water, sterilized Kanawha River water, agreement with accepted values and lend credibility to the and Kanawha River water as is, each containing about 65 laboratory procedures used and other results achieved (21). mg/L ethylene oxide were tested at 25 "C and pH 7.4, the The results of hydrolysis tests using 1.0 and 3.0 wt % NaCl Environ. Sci. Technol., Vol. 17, No. 2, 1983
109
Table 111. Comparison of Calculated and Experimental Hydrolysis/Hydrochlorination Results log k: t,,,, h NaCl, %
exptl
0 1 3 a log k obtained Reference 4: see
-6.21 ( 3 1 4 ) -6.14 (265) -6.06 (224) from k in s-l. also Table IV.
chlorohydriniglycol ratio prior lit. calcC
model fitb
prior model fitb lit. calcC
exptl
-6.17 (285) - 5.99 (188) 0.11 0.09 0.8 -6.10 (242) -5.71 (99) 0.23 0.28 2.4 Calculated by using experimental data and equations cited in text; see also Table IV.
Table IV. Hydrolysis/Hydrochlorination Rate Constants for Ethylene Oxide a t 25 "C and pH 7a k, 0.57 X lO-'s-' k, 9 X lo-' L M-' k, 1x L M-l CA= ~ 0.57 M = 3% NaCl
\k=-
CALCULATED R E L A T I O N S H I P S
\\
k,i k,i
0.24 3.67
X X
lo-' L M-' lo-, L M-'
s - ~ ~ , ~ s - ~
a Reference 4. Estimated by calculation by Bogyo e t al. ( 4 ) with a ratio suggested by Bronsten e t al. ( 2 3 ) . Determined value in work reported herein is 0.305 X L M-I s-l from eq 7.
Table V. Biodegradation biooxidation, % o f theoretical oxygen demand 2
4
6
8
1
0
1
2
material tested
day 5
day 10
day 15
day 20
ethylene oxidea ethylene glycolb2c ethylenechlorohydrind glyoxal
5c 3gC 0 7
22 73 13 65
40
52 96 57 76
PH
Flgure 3. Ethylene oxide hydrolysis, authors' results vs. literature data (ref 4).
indicated pseudo-first-order reactions with log k's of -6.14 and -6.06 s-l, respectively. Half-lives associated with these were 265 and 224 h. During the hydrolysis tests with these materials, the loss of ethylene oxide and the formation of ethylene chlorohydrin were determined. This permitted the generation of chlorohydrin/glycol ratios in Table 111. These were 0.11 and 0.23 for 1% and 3% NaCl solutions, respectively. The chlorohydrin/glycol ratio was shown by Bogyo et al. (4) to be determined by the equation chlorohydrin/glycol =
[k3i + k4iCH30+1CAi kl
+ k2CHsO++ k3C0H-
(6)
where kl, k2, and k3 are the specific rate constants for noncatalytic, acid-catalyzed, and base-catalyzed reactions, respectively, and C, k3i, k4irefer to the concentration and specific rate constants for the anion. Substitution of the appropriate value for each of the terms at pH 7 and 25 OC (Table IV) results in a ratio of 2.4 as shown in the Bogyo article ( 4 ) and indicates that the ratio may be simplified to ratio = (k3j/k,)CAi (7) The rate constant and half-life in marine water may be calculated from = [ki + ( ~ ~ ( C A I ) ] (8) t l l , = In 2/12 = O,693/k (9) Substitution of values from Table IV results in k equal to 0.194 x s-l. This was used to calculate the published half-life of 99 h, as indicated by Bogyo et al. ( 4 ) of the Syracuse Research Corp. Our determined chlorohydrin/glycol ratios were each used to calculate a k3i by using eq 7. The results (0.36 X 110 Environ. Sci. Technol., Vol. 17,
No. 2,
1983
85 46 74
a Mean values of eight tests; standard deviations were 4, 12, 16, and 1 6 for the four periods; these are unusually high-presumably due to the concurrent hydrolysis and perhaps other reactivity of the material, Mean values of four tests; standard deviations were 16, 13, 7, and 6. Shell investigators reported 3% for ethylene oxide and 36% for ethylene glycol at 5 days ( 2 4 , 2 5 ) . Lamb and Jenkins report 12, 52%, 71%, and 78% for ethylene glycol for days 5, 10, 15, 20, respectively (25). Lamb and Jenkins report OR, 16%, 74%, and 87% for days 5, 10, 15, 20, respectively ( 2 5 ) .
lo4 and 0.25 X lo4 L M-l s-l) were averaged to produce the single k3ivalue of 0.305 X lo4 L M-' s-', which is about one-tenth that calculated by SRC. With this average value, ratios were calculated by using eq 7 and reaction rate constants were calculated by using eq 8 and half-lives by using eq 9. Calculated and experimental data are compared with SRC data in Table 111. The results show that use of the averaged jZgi value permits calculation of ratios, rate constants, and half-lives reasonably representative of those found experimentally. Biooxidation. Data presented in Table V for EO and some possible metabolites show that these materials biodegrade at reasonable rates in a lightly seeded dilution-bottle test; biooxidation of EO was 52% in 20 days. In a biological waste treatment system with a high concentration of adapted microorganisms, biooxidation would be much faster, perhaps a matter of hours. Miller has published a study of glycols (26) similar to that done by SRC for EPA on epoxides. Degradative pathways are presented along with additional rate information on river water and other systems. Degradation was relatively rapid, and persistence in the environment would be precluded. There was no evidence to suggest that glycols would bioaccumulate.
Table VI. Aquatic Toxicity Data on Ethylene Oxide and Related Materials toxicity, LC,, (95% conf limits), mg/La material tested
test procedure
ethylene oxide range-finding, static, aerated range-finding, static, sealed under oxygen definitive static acute (no aeration) static acute static acute ethylene glycol definitive static acute static acute static acute ethylene range-finding. - , static chlorohydrin acute definitive static acute static acute static acute glyoxal definitive static acute sodium lauryl* definitive static acute sulfate static acute static acute
test organism
24 h
fathead minnow
274 (150-500)
fathead minnow
8 6 (50-150)
fathead minnow
90 (63-125)
Daphnia magna
>300 27 0 260 brine shrimp > 500 350 570 fathead minnow > 1 0 000 Daphnia magna > 1 0 000 > 20 000 brine shrimp fathead minnow > 500 fathead minnow
Daphnia magna brine shrimp fathead minnow fathead minnow Daphnia magna brine shrimp
48 h
96 h
8 9 63-125) 300 137 83-179) 200 150-243) > 500 1000 490 > 1 0 000 > 10 000
995 768 67 5 >loo0 550 7.4 (6.1-8.9) 8.5 (6.8-10.5) 10 12 7.2 13
84 (73-96)
> 10 000
164 (120-236) 163 (140-185) 100 (50-200) 680 230 6.6 (5.8-7.5) 7.3 (5.8-9.3) 5.6 (3.3-8.2) 4.8 (3.1-6.5) 4.6 (2.8-6.4) 1.5
67 (49-84) 112 (90-131) 215 6.6 (5.8-7.5) 6.9 (5.3-90)
Measured by procedures published by EPA and ASTM ( 7 ) : some test modifications were required t o meet sample size Sodium lauryl sulfate is limitations and dissolved oxygen requirements. LC,, refers to the median lethal concentration, included as a standard for a check on procedures and health of test organisms. a
Toxicity to Aquatic Life. The toxicity of EO and related materials to aquatic life is shown in Table VI. The 96-h LC,, (the dosed concentration that killed half of the test organisms in 96 h) of EO to fathead minnows was determined to be 84 mg/L; a literature value for goldfish was 90 mg/L at 24 h (25). Testing with a fish-food organism (Daphnia magna) showed a mean LC50 of 212 mg/L at 48 h; with brine shrimp the mean LC,, of EO at 48 h was 745 mg/L. Upon hydrolysis to ethylene glycol (EG), toxic levels rise to above 10000 mg/L (1%). If reacted to form ethylene chlorohydrin (EC), the 96-h LCw to fathead minnows is about 90 mg/L. As discussed previously, none of these materials is persistent; between biochemical oxidation, reactivity, volatilization, and dilution, the reaching of toxic levels except in a very localized incident area would seem unlikely. Effect on Treatment Plant Biomass. The adverse effect level of ethylene oxide on activated sludge microorganisms or the ICho(concentration that inhibited growth 50%) was determined to be in the range of 10-100 mg/L. The IC,, of ethylene glycol was above 10OOO mg/L. Risk Analysis. The foregoing information can be used to establish the probability and severity of risk at an actual site, a hypothetical or "typical" site, or a broad area. First, an expected environmental concentraiton (EEC) is calculated. The EEC then is compared with toxicity values, by taking into account expected exposure periods. Risk analysis techniques have been summarized in texts by Neely (27) and Conway (28). An overview of the data indicates that due to biooxidation, readivity, volatilization, and dilution, the reaching of toxic levels except in very localized incidents would be unlikely. Conclusions (1)The desorption rate of EO from natural waters is about 0.36 times that of oxygen under the same conditions.
*
Values were estimated for two situations. (2) The hydrolysis half-life of EO is about 14 days in fresh water at 25 "C. Temperature effects are greater than pH effects at pH levels usually encountered. (3) The hydrolysis/hydrochlorination half-life of EO in salt water is about 9 days at 25 OC. The ratio of chlorohydrin to glycol formed is about 0.2 at 3% NaC1; the ratio is directly proportional to salt concentration. (4) The aquatic toxicity of EO was found to be moderate, e.g., 96-h LC50 with fathead minnows of 84 mg/L, 48-h LCw with Daphnia of 212 mg/L, and 48-h LCw with brine shrimp of 745 mg/L. (5) Biooxidation of EO, EG, and EC proceeds at rates indicating nonpersistence. In a lightly seeded BOD test, 52% of EO was oxidized in 20 days without prior acclimation. (6) This information can be used to evaluate environmental risks at specific sites. However, due to biooxidation, reactivity, volatilization, and dilution, the reaching of toxic levels except in very localized incidents would seem unlikely.
Acknowledgments The advice and guidance of G. M. Alsop, L. J. Priestley, W. F. Tully, and B. E. Wilkes regarding gas chromatography equipment and operation were invaluable. Registry No. EO, 75-21-8; ethylene glycol, 107-21-1;ethylene chlorohydrin, 107-07-3; glyoxal, 107-22-2.
Literature Cited (1) Chemical Regulation Reporter 1980, Aug 1, 534-5. (2) Chemical Regulation Reporter 1979, Oct 10, 1358-60. (3) Chemical Regulation Reporter 1980, Mar 7, 1883. (4) Bogyo, D. A.; et al. "Investigation of Selected Environmental Contaminants: Epoxides"; EPA-560/11-80-005, Office of Toxic Substances, U.S. Environmental Protection Agency, Envlron. Sci. Technol., Vol. 17, No. 2, 1983
111
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Washington, DC 20460, Nov 1980. “Standard Methods for the Examination of Water and Wastewater”, 14th ed.; American Public Health Association: Washington, DC, 1975. Price, K. S.; Waggy, G. T.; Conway, R. A. J. Water Pollut. Control Fed. 1974, 46,63-77. Committee on Methods for Toxicity Tests with Aquatic Organisms, “Methods for Acute Toxicity Tests with Fish, Macroinvertebrates and Amphibians”; EPA-660/3-75-009, Apr 1975. Alsop, G. M.; Waggy,G. T.; Conway, R. A. J. Water Pollut. Control Fed. 1980,52, 2452-2456. Liss, P. S.; Slater, P. G. Nature (London) 1974, 247, 181-184. Mackay, D.; Leionen, P. J. Environ. Sci. Technol. 1975,9, 1180. Neely, W. B. Proceedings of the 1976 National Conference of Hazardous Material Spills, 1976, pp 197-200. Dilling, W. L. Enuiron. Sci. Technol. 1977, 11, 405-409. Sweeris, S. Prog. Water Technol. 1979,11, 37-47. Smith, J. H.; Bomberger, D. C. “Hydrocarbons and Halogenated Hydrocarbons in the Aquatic Environment“; Afghan, B. K., MacKay, D., Ed.; Plenum Press: New York, 1980, pp 445-451. Smith, J. H.; e t al. Enuiron. Sci. Technol. 1980, 14, 1332-1337. Smith, J. H.; et al. Chemosphere 1981, 10, 281-289. Cawse, J. N.; et al. In “Kirk-Othmer Encyclopedia of Chemical Technology”, 3rd ed.; Wiley: New York, 1980; Vol. 9, pp 432-471.
(18) MacCormack, K. E.; J. H. B. Chenier Ind Eng Chem. 1955, 47, 1454-1458. (19) Reid, R. C.; e t al. ”The Properties of Gases and Liquids”, 3rd ed.; McGraw-Hill: New York, 1977; p 567. (20) Rathburn, R. E.; Tai, D. Y. Water Res. 1981,15,243-250. (21) Mabey, W.; Mill, T. J. Phys. Chem., Ref. Data 1978, 7, 383-415. (22) Mill, T.; Mabey, W. R.; Hendry, D. G. “The Fate of Selected Pollutants in Freshwater Aquatic Systems. Protocol 1: Hydrolysis”; SRI International: Menlo Park, CA 94025, Contract 68-03-2227, 1978. (23) Bronsted, J. N.; Kilpatrick, M.; Kilpatrick, M. J. Am. Chem. SOC.1929,51, 428. (24) Bridie, A. L.; Wolff, C. J. M.; Winter, M. Water Res. 1979, 13,623-626. (25) Verschueren, K. “Handbook of Environmental Data on Organic Chemicals”; Van Nostrand Reinhold: New York, 1977. (26) Miller, I. M. “Investigation of Selected Potential Environmental Contaminants: Ethylene Glycol, Propylene Glycols, and Butylene Glycols”;Franklin Research Center: Philadelphia, PA, 1979 (prepared for U.S. EPA, EPA 560/ 11-79-006). (27) Neely, W. B. ”Chemicals in the Environmenti Distribution, Transport, Fate, Analysis”; Marcel Dekker: New York, 1980. (28) Conway, R. A. “Environmental Risk Analysis for Chemicals”; Van Nostrand Reinhold: New York, 1982.
Received for review April 9, 1982. Accepted October 4, 1982.
Smog Chamber Studies of NO, Transformation Rate and Nitrate/Precursor Relationships Chester W. Spicer Battelle Columbus Laboratories, Columbus, Ohio 43201 An environmental chamber study of nitrogen oxide reactions is described. The aim of the investigation was to determine (1)the rate of NO, transformation to nitrate products, (2) the environmental factors that influence the transformation rate, (3) the relationship between nitrate products and their hydrocarbon and NO, precursors, and (4) relationships among NO,, 03,and nitrates. The experiments made use of a 17-component hydrocarbon mixture designed to represent the major organic constituents of urban air. The results of the chamber experiments demonstrate that the transformation of NO2 to nitrate products follows pseudo-first-order kinetics, and the transformation rate constant is proportional to the initial NMHC/NO, ratio. Nitrogen balances of 290% were obtained when “OB deposition to the chamber surface was taken into account. The fractional conversion of NO, to products was shown to depend on initial NMHC/NO,. For this particular hydrocarbon mixture, the PAN/HN03 ratio was directly proportional to initial NMHC/NO, for NHMC/NO, 120. For a typical urban NMHC/NO, ratio of 8, about 3 times as much HNO, as PAN is produced in the chamber. The implications and limitations of the experimental results are discussed. Introduction The conversion of nitrogen oxides (NO, = NO + N0.J to gaseous and particulate nitrates in the atmosphere is an important but poorly understood process. The rate of the conversion affects ozone formation and the ultimate fate of the nitrogen oxides. The nitrate products that are formed may play a role in the acidification of precipitation 112
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and have been reported to have a deleterious effect on respiratory function and on human health in general. In view of the vital role of these species in atmospheric chemistry, much more needs to be learned about the mechanisms and the dynamics of the processes that convert NO, to nitrates. The conversion of NO, to products has been the subject of several previous laboratory studies (1-8). However, little has been published on the rate of transformation of NO, to nitrates and the factors that affect this rate. There is also limited information available concerning the relationships between the nitrate reaction products and the hydrocarbon and nitrogen oxide precursors. Since the environmental effects of the major reaction products nitric acid, peroxyacetyl nitrate (PAN), and particulate nitrate are different, it is important to understand how the distribution of these products varies with changing precursor concentrations and environmental conditions. It is also necessary to understand these relationships so that we may predict the consequent changes in NO, reaction product distribution of (1) proposed pollutant control strategies, (2) changes in hydrocarbon or NO, emissions due to future energy or environmental constraints, (3) changes in hydrocarbon composition due to changing energy sources or transportation modes, or (4) other future changes in environmental conditions. This study was undertaken to provide a better understanding of atmospheric NO, transformation rate, the factors that affect the rate in urban air, and the relationship between nitrate products and their hydrocarbon and NO, precursors. Variations in the initial hydrocarbon and NO, concentrations and ratios were employed to derive
00 13-936X/83/0917-0112$01.50/0
0 1983 American Chemical Society
