system. Although a segmental control system may be a practical necessity, external costs of various pollutants, in addition to direct operating costs, must be imputed in evaluating control strategies, schemes, and processes. APPENDIX
Brief Description of the Five S t a c k Gas Desulfurization Processes All process data represent the state of technology in late 1973. Limestone S l u r r y Scrubbing. Stack gas is washed with a recirculating slurry (pH of 5.8-6.4) of limestone and reacted calcium salts in water using a two-stage scrubber system for particulates and SO2 removal. Limestone feed is wet ground prior to addition to the scrubber effluent hold tank. Calcium sulfite and sulfate salts are withdrawn to a disposal area for discard. Lime S l u r r y Scrubbing. Stack gas is washed with a recirculating slurry (pH of 6.0-8.0) of calcined limestone (lime) and reacted calcium salts in water using a two-stage venturi scrubbing. Lime is purchased from “across the fence” calcination operation, slaked, and added to both circulation streams. Calcium sulfite and sulfate are withdrawn to a disposal area for discard. Magnesia S l u r r y Scrubbing-Regeneration to HzS04. Stack gas is washed using a two-stage venturi scrubbing. Water is utilized for removal of particulates, and a recirculating slurry (pH 7.5-8.5) of magnesia (MgO) is utilized for removal of S 0 2 . Makeup magnesia is slaked and added to cover only handling losses since sulfates formed are reduced during regeneration. Slurry from the SO2 scrubber is dewatered, dried, calcined, and recycled, during which concentrated SO2 is evolved to a contact sulfuric acid plant producing 98% acid. Sodium Solution Scrubbing-SO2 Reduction to Sulfur. Stack gas is washed with water in a venturi scrubber for removal of particulates and then washed in a valve tray scrubber with a recirculating solution of sodium salts in water for SO2 removal. Makeup sodium carbonate is added to cover losses due to handling and oxidation of sodium sulfite to sulfate. Sodium sulfate crystals are purged from
the system, dried, and sold. Water is evaporated from the scrubbing solution to crystallize and thermally decompose sodium bisulfite, driving off concentrated SO1. The resulting sodium sulfite is recycled to the scrubber, and the SO2 is reacted with methane for reduction to elemental sulfur. Catalytic Oxidation. Stack gas is first cleaned of particulates by a high-temperature electrostatic precipitator. Then the SO2 is catalytically converted to SOs, and available excess heat is recovered. The SO3 reacts with moisture in the stack gas to form H2S04 mist which is scrubbed in a packed tower using a recirculating acid stream to yield 80% acid. The mist is removed by a Brink mist eliminator, and the clean 254’F gas is exhausted to the stack.
Acknowledgment I am indebted to T. Y. Yan of Mobil Research and Development Corp. for his help in constructing the material balance tables and A. M. Chung of Drexel University for his valuable suggestions.
Literature Cited (1) Leontief, W., Reo. Econ. Stat., 52,262-71 (Aug. 1970). (2) Ayres, R. U., Kneese, A. V., Am. Econ. Reu., 59,282-97 (1969).
(3) Kahn, R. E., Proc. 1970 Soc.-Stat. Sec., Am. Stat. Assoc., pp 207-14. (4) Russel, C. S., Am. Econ. Reu., 63,236-43 (1973). (5) McGlarmery, G. C., Torstrick, R. L., “Cost Comparison of Flue Gas Desulfurization Systems,” Fuel Gas Desulfurization Symposium, Atlanta, Ga., Nov. 1974. (6) Stanford Research Institute, Chemical Economics Handbook, Menlo Park, Calif., 1974. (7) Sulfur Oxide Control Technology Assessment Panel on Projected Utilization of Stack Gas Cleaning Systems by SteamElectric Plants, “Final Report,” ASTD 1569, April 1973. (8) Jenkins, R. E., McCutchen, G. D., Environ. Sci. Technol., 6, 884-8 (1972). (9) Environmental Protection Agency, “The Economics of Clean Air,” Annual Report of the Administrator of the EPA to the Congress of the U.S., March 1972. (10) Faith, W. I., Keyes, D. B., Clark, R. L., “Industrial Chemicals”, 3rd ed., John Wiley, New York, N.Y., 1965. (11) Kirk, R. E., Othmer, D. F., Encyclopedia of Chemical Technology, 2nd ed., 1964, ibid. Received for review January 29, 1975. Accepted August 27, 1975. This work was supported by a grant from Drerel C’niuersity.
Atmospheric Oxidation of Chlorinated Ethylenes Bruce W. Gay, Jr.*, Philip L. Hanst, Joseph J. Bufalini, and Richard C. Noonan Environmental Protection Agency, Environmental Science Research Laboratory, Atmospheric Chemistry and Physics Division, Research Triangle Park. N.C. 277 11
Chlorinated ethylenes are released into the air in great quantities. The US.Tariff Commission gives the following 1974 U.S. production figures: vinyl chloride, 2.5 X lo9 kg; trichloroethylene, 2.0 X lo8 kg, and tetrachloroethylene, 3.3 X lo8 kg. The vinyl chloride monomer produced is largely converted to polyvinyl chloride, but it has been estimated that approximately 6% of the monomer is lost to the air during processing, corresponding to 1.5 X lo8 kg in 1974. The 1974 production of all halogenated organic compounds in the United States was more than five billion kilograms (11.
These halogenated organic compounds do not occur naturally in the environment, and therefore much attention is being given to their effects on human health and welfare. 58
Environmental Science 8 Technology
The fully halogenated compounds, such as trichlorofluoromethane and dichlorodifluoromethane, are chemically unreactive and have not been implicated in any direct effects on human health. Their lack of reactivity, however, allows them to remain in the atmosphere long enough to reach the stratosphere, where a detrimental interaction with stratospheric ozone may be taking place (2, 3 ) . The halogenated ethylenes seem to be sufficiently reactive for them to be fully photooxidized in the troposphere. While this reactivity may be a desirable property from the point of view of stratospheric effects, it apparently is associated with direct effects on human health. Vinyl chloride has recently received notoriety by being linked to angiosarcoma, a rare type of cancer in the human liver. Exposures
Ethylene and chlorinated ethylenes were photooxidized in air in the presence of nitrogen dioxide with ultraviolet light. Analyses were performed by long-path infrared spectroscopy, wet chemical, and chemiluminescent procedures. Reactivities of the ethylene compounds fall in this decreasing order: 1,l-dichloroethylene > 1,2-dichloroethylene-trichloroethylene > ethylene > vinyl chloride >> tetrachloroethylene. Reaction sequence involves photolysis of nitrogen dioxide, formation of ozone, conversion of the olefinic
compounds to a highly oxygenated intermediate, attack by the intermediate on olefin and nitric oxide molecules, formation of transitory epoxides, and rearrangement of epoxides through chlorine atom movement to yield chlorinated acetaldehyde or chlorinated acetyl chloride. The chlorinated aldehydes are further oxidized to more stable final products, including a small yield of chlorinated peroxyacetyl nitrates.
to vinyl chloride have been greatest in occupational situations and when vinyl chloride has been used as the propellant in home use products, such as in hair spray and insecticides ( 4 ) . Vinylidene chloride (1,l-dichloroethylene) has been shown to attack the liver in test animals ( 5 ) .Trichloroethylene, the widely used dry cleaning solvent and degreasing agent, also shows evidence of the same type of effects. I t is thus clear that information on the atmospheric degradation of halogenated compounds is critically needed. Of primary interest are the rates of photooxidation and the identity of photooxidation intermediates and final products. Previous laboratory research work has yielded insights into a number of aspects of the photooxidation processes for the halogenated ethylenes. Valuable studies have been carried out, for example, by J. Heickien and his colleagues working a t the Pennsylvania State University under an Environmental Protection Agency research grant (6). Nevertheless, all processes and products have not been identified or understood, and further measurements were undertaken in the present project.
injection ports to introduce samples into the cell were used. The chamber ends were capped with flat Plexiglas end plates 3.18 cm thick, using Teflon gaskets. The internal volume of the chamber was 690 1. A large displacement vacuum pump was connected to one end plate with a ball vacuum valve through which the chamber could be evacuated to a pressure 150
66
2.07
1.75
2.90
None
2.66
115
66
2.07
0.80
1.45
0.47
2.83
> 180
7
0.07
0.27
0.42
0.12
190
1.77 4.63 4.85
2.21 3.0 2.15
> 160
5.0
3.0
3.45 5.0
1000
1400
1800
2200
2600
3000
Figure 2. Spectra of ethylene and NO2 Top spectrum of ethylene (2 ppm) and NOn (1 ppm) before irradiation. Lower spectrum after 120 min of irradiation [kdNoz)= 0.6 min-', p = 760 torr, 500-m path1
60
Environmental Science & Technology
diation. For graphic display of spectra, the resolution of initial spectral data was degraded by a factor of four to give spectra of four wave number resolution. Figure 3 shows growth and decay of reactants and products of the photooxidation of ethylene and NOz. After 240 min of irradiation, approximately 82% of the ethylene had reacted. The maximum ozone concentration of 1.23 ppm was formed a t 190 min. After 160 min of irradiation, the formaldehyde concentration became constant with increasing irradiation time. Vinyl Chloride. Spectral data of vinyl chloride-NO2 photooxidation are shown in Figure 4. An imbalance of CO2 between the sample and reference air causes the inversion of the COz band. The concentration-time plots for the irradiation of vinyl chloride in the presence of NO:! are shown in Figure 5. At 160 min of irradiation time approximately 40% of the vinyl chloride had been reacted. The major products observed using in situ infrared were: formic acid, hydrochloric acid, carbon monoxide, formaldehyde, and ozone. Trace amounts of formyl chloride and nitric acid were also detected but not quantified. The formyl chloride is thermally unstable and decomposes to HCl and CO. A small fraction of a ppm of COS is probably produced but this goes undetect-
ed due to spectral interferences mentioned earlier. A t the conclusion of the irradiation (160 min), 76% of the reacted carbon could be accounted for as products and 87% of the chlorine as hydrochloric acid in the gas phase. Weakly absorbing unidentified peaks were observed a t 790 cm-', 1165 cm-I, and 1292 cm-' in the spectra, Figure 4. In the photooxidation of ethylene and NO2 (Figure 2), these absorbing bands are absent. Since peroxyacetyl nitrate has similar absorption bands a t 792 cm-l, 1162 cm-', and 1300 cm-l, these unknown bands may be those of a chlorinated peroxyacetyl nitrate. In an attempt to identify these bands, 3-chloropropene and NO2 were photolyzed to produce chlorinated peroxyacetyl nitrate. The chlorinated peroxyacetyl nitrate produced had a stronger absorption a t 790 cm-' than 1165 cm-I when com-
4 63
I
I
40
60
I
I
1
I
I
eo
100
120
140
160
-I
2.0
1.6
-
12
.e
U
3
4
n
0
30
60
90
120
150
TIME
180
210
240
IN f m n l
0
Figure 3. Reactants and products of ethylene and NO2 photooxidation
20
TIME I N m n
Figure 5. Photooxidation products of vinyl chloride and NO2
-4 v1
1000
1400
1800
22.00
2600
3000
Figure 4. Spectra of vinyl chloride and NO2 Top spectrum of vinyl chloride ( 5 ppm) and NO2 (1.5 ppm) before irradiation. Lower spectrum after 100 rnin of irradiation. [kflNOz) = 0.6 min-', p = 760 torr, 500-m path)
Volume 10, Number 1. January 1976
61
600 CM-l
1400
1000
1800
2200
2600
3000
Figure 6. Spectra of 1,l-dichloroethylene and NO2 Upper spectrum of l,ldichloroethyle% (5 ppm) and NO1 (1 ppm). Lower spectrum after 100 min of irradiation. [kdNo2)= 0.6 min-', p = 760 torr, 500-m path]
\
"0
20
40
CHLOROACETYL CHLORIUE
60
80 1W TIME IN Imml
120
140
160
Figure 7. Reactants and products of 1,I-dichloroethylene and NO2 irradiation
pared against peroxyacetyl nitrate. In the photooxidation of vinyl chloride and NO2 the absorption peaks and relative absorption strengths were similar to that of chlorinated peroxyacetyl nitrate. Reaction mechanism discussions will show that this compound is a logical reaction product. 1,l-Dichloroethylene. The 1,l-dichloroethylene photooxidized much more rapidly than vinyl chloride. The products detected were those observed for vinyl chloride photolysis plus phosgene and chloroacetyl chloride (Figure 6). Chloroacetyl chloride has also been reported as a product from the ozonolysis of 1,l-dichloroethylene (15).After 140 min of irradiation, approximately 83% of the compound had reacted (Figure 7). 62
Environmental Science & Technology
1,2-Dichloroethylene. The photooxidation of 1,2-dichloroethylene is shown in Figures 8 and 9. The reactivity of this halocarbon is greater than vinyl chloride but less than 1,l-dichloroethylene. The major products of photooxidation are similar to those observed for vinyl chloride, except that no formaldehyde is seen. Unknown absorbances a t 791 cm-l and 1165 cm-l might be that of a dichlorinated PAN. No experiments were included to verify this but it is assumed to be likely in view of the formation of chlorinated PAN from vinyl chloride. The gas phase concentration of hydrochloric acid accounts for 45% of the reacted 1,2-dichloroethylene. Hydrochloric acid absorbed on the chamber wall and gas phase formyl chloride should account for that chlorine not otherwise present in the gas phase. Trichloroethylene. The photooxidation of trichloroethylene is shown in Figures 10 and 11. The spectral data show the phosgene absorption a t 850 cm-' and 1820 cm-l, the formyl chloride in the 710-720-~m-~region is overshadowed by the strong absorptions of dichloroacetyl chloride, a major product. Tetrachloroethylene. The photooxidation of tetrachloroethylene completes the study of the reactions of the substituted ethylenes in the LPIR chamber. I t is the least reactive of the chlorinated ethylenes. Only 0.37 ppm or 7% of the compound reacted over a period of 3 hr. The spectral data in Figure 12 show phosgene, formic acid, carbon monoxide, hydrochloric acid, and trichloroacetyl chloride. As can be seen from Figure 13, the reactivity of this compound is quite low. Only a small concentration of ozone was produced in a three hour irradiation. Vinyl chloride-Ozone Reaction. Our photooxidation of vinyl chloride in the presence of NO2 showed a reactivity approximately 60% that of ethylene. I t was expected therefore that the ozone-vinyl chloride reaction would be correspondingly slower than the ozone-ethylene reaction. The results of the reaction of ozone and vinyl chloride are plotted in Figure 14. The primary products of the reaction are carbon monoxide, formaldehyde, formic acid, and a small quantity of hydrochloric acid. The rate constant for the reaction, assuming a second-order mechanism, is 0.34 X pprn-l min-'. This value is approximately y10 that of the ozone-ethylene reaction (3 X ppm-l min-') recommended by Garvin and Hampson (16). The ratio
60C Q4-l
1000
1400
1800
2200
2600
3000
Figure 8. Spectra of 1,2-dichioroethylene and NO2 Upper spectrum of 1,2-dichloroethylene(5 ppm) and NOn (1 ppm) before irradiation. Lower spectrum after 100 rnin of irradiation. [ k d ~ ~=20.6 1 min-', p = 760 torr, 360 m]
divided by kethylene is much greater than that found by Williamson and CvetanoviE ( 1 7 ) in carbon tetrachloride solution. The stoichiometry of the reaction from Figure 14 suggests that more ozone than vinyl chloride is consumed. The ozone concentration curve had not been corrected for thermal and heterogenous degradation of ozone. The ozone degraded in the LPIR chamber a t a rate of 4.7% per hour. If this correction is used for the ozonolysis run, the amount of ozone consumed in 3 hr is only approximately 0.5 ppm. Thus, more vinyl chloride than ozone is consumed during the reaction-an observation compatible with other ozoneolefin reactions (17, 1 8 ) . 335-Ft3 Chamber. Ethylene-NO. The photolysis of ethylene in the presence of NO, (>go% NO) in the 335-ft3 chamber is shown in Figure 15. Data from this chamber are shown primarily for comparative purposes. After 300 rnin of irradiation, approximately 67% of the ethylene reacted. The NO2 maximum occurs a t approximately 90 min, and ozone reaches a maximum of 0.82 ppm a t 300 min. The formaldehyde maximum of 1.06 ppm also occurred a t 300 min. When the NO, is mainly NO, an induction period is necessary for the conversion of NO to NO2 before ozone forms. With NO2 as the reactant ozone starts to form immediately, Figure 3. Vinyl Chloride-NO. The reaction of vinyl chloride-NO in the 335-ft3 chamber is shown in Figure 16. The conditions for photooxidation are the same as those of ethyleneNO. After 5 hr of irradiation 39% of the vinyl chloride reacted. The NO2 maximum occurs a t 130 min. The ozone observed after 330 rnin was 0.44 ppm. The concentration of formaldehyde reached a maximum of 0.32 pprn a t 140 rnin and leveled off a t this value for the remainder of the run. This suggests that the production rate of formaldehyde is approximately equal t o the rate of destruction after 140 min. In terms of NO oxidation, vinyl chloride is 70% as reactive as ethylene. In terms of compound reacted, vinyl chloride is 60% as reactive as ethylene. The ozone produced a t the close of the run, 330 min, is one half as great for the vinyl chloride case compared to the ethylene case. However, for vinyl chloride photooxidation, the slope of the kvinyl chloride
1-
00
20
40
60
80
100
120
140
160
TIME IN ( m n l
Figure 9. Photooxidation of 1.2-dichloroethyleneand NO2
ozone-time curve is still positive, even a t the close of the run. This suggests that with continued irradiation, the ozone concentration would increase. At the close of the irradiation, 20% of the reacted vinyl chloride appeared as formaldehyde, whereas 33% of the reacted ethylene appeared as formaldehyde in ethylene photooxidation. The high formaldehyde yield from ethylene is expected since both ends of the molecule can give rise to formaldehyde. Discussion
Ethylene. The photooxidation of ethylene in air through the action of NO2 and sunlight is a complex reaction that Volume
IO, Number 1, January 1976
63
Figure 10. Spectra of trichloroethylene and NOP Upper spectrum of trichloroethylene (5 ppm) and NO? (1 ppm) before irradiation Lower spectrum after 100 min of irradiation [kdNOz)= 0 6 min-’, p = 760 torr, 500-m path)
has been the subject of much previous study (19). Attacking species in this case include, 0, 0 3 , OH, and OzH. Principal reactions include:
hydrogen in the molecule, the list of stable reaction products is extended to include hydrogen chloride and chloroperoxyacetyl nitrate. The hydrogen chloride results from the decomposition of formyl chloride, which is seen in the spectrum as an intermediate compound. The addition of ozone to the double bond can go either of two ways. In one speculative configuration, formaldehyde, Con, and HCl can be produced:
Further reaction of formaldehyde yields formic acid, hydrogen peroxide, carbon monoxide, carbon dioxide, and water. Nitrogen dioxide is continually restored by the fast reaction between NO and ozone:
so
-
+0
+
KO.
0.
HC1
+
COL
+
H-CO
(10)
In the other configuration, formyl chloride, CO, and H20 can be produced:
(5)
However, HOz radicals, HC03 radicals, and HzCOO diradicals compete with the ozone in the oxidation of NO: SO
+
RO,
-
NO1
+
RO
(6)
This allows excess ozone to build up in the system. No nitrogen-containing organic products have ever been detected in the ethylene-NO2 reaction. The nitrogen oxides are removed from the system by reaction with ozone, leading to nitric acid: NO.
+
-- + -
0,
3h TRICHLOROETHYLENE
KO
0-
(i)
NO
+
KO-
N,O,
(8)
N-0;
+
HO
2HS0,
(9)
DICHLOROACETVL CHLORIDE
-E
E
;
2
8
1
0 0
The nitric acid appears in the vapor phase, in the fine particles, and on the vessel walls (20). Vinyl Chloride. When a chlorine atom has replaced a 64
Environmental Science & Technology
I
20
40
M
Bo
100 TIME
120
140
160
180
2W
IN lminl
Figure 11. Reactants and products of trichloroethylene and NO2 irradiation
Figure 12. Spectra of tetrachloroethylene and NO2 'Upper spectrum of tetrachloroethylene (5 ppm) and NO2 (1 ppm). Lower spectrum after 100 rnin of irradiation. [ k q ~=~0.6 ) min-', p = 760 torr, 500-m path]
(2) Attack of RO2 radicals (including diradicals) on the double bond:
HClCO
f
H?O
+ CO
(11)
The formyl chloride decomposes thermally to CO and HC1. Hisatsune and Heicklen (21) report the half-life of HClCO a t room temperature to be about 20 min. The diradicals, H2COO and HClC00, will also exist as intermediates prior to the formation of the final products. The buildup of ozone in the reaction indicates that the diradicals or other types of peroxide radicals, RO2, must react with the NO to yield NO2: RO? f SO
-
RO
+
YO.
When the monochloroethylene oxide rearranges, monochloroacetaldehyde is formed:
(12)
Such oxygen donation by a diradical is energetically strongly favored because after oxygen loss, the remainder of the diradical is a thermodynamically stable carbonyl compound. The diradicals can even oxidize SO2 to SO,?, as shown by McNelis ( 2 2 ) .The highly oxygenated intermediates shown above in square brackets may also be considered to be among the attacking species, similar to RO2 radicals. A high degree of free radical attack is required to account for the products. The production of between 2 and 3 ppm of ozone, for example, requires that 2 or 3 ppm of NO have been oxidized to NO2 through attack of RO2 radicals. Ozone consumed in reactions with olefin and NO2 also had to be produced by the attack of RO2 radicals on the NO with subsequent photodissociation of the NO2. The production of chlorinated peroxyacetyl nitrate requires the prior formation of chloroacetaldehyde. This is best accounted for by a rearrangement of an unstable monochloro ethylene oxide. This oxide may be formed in several ways: (1) Addition of oxygen atom to the double bond: 0 0
20
40
60
60 100 TIME IN h n )
120
140
160
180
Figure 13. Reactants and products of tetrachloroethylene-NO2 photooxidation Volume
IO, Number 1, January 1976 65
The chlorinated aldehyde proceeds to the corresponding PAN-type compound through the following mechanism (24):
L
H
H
J
TIME iHni
Figure 14. Ozonolysis of vinyl chloride
L
H
J
Monochloroperoxyacetyl Nitrate
TIME $Nl n l n l
Figure 15. Ethylene and NO photooxidation 355-ft3chamber 20
I
1
1
1
1
I
I
I
I
1
I
)
Figure 16. Vinyl chloride and NO photooxidation-355-ft3 chamber
H
Hull e t al. (15) have invoked this same rearrangement. They call the intermediate a “hot” epoxide. Van Duuren (23) has postulated the existence of transitory epoxides as activated carcinogenic intermediates during metabolism of chlorinated olefins. In this vinyl chloride case, either the transfer of hydrogen atom or the transfer of the chlorine atom from one carbon to the other would yield the chloroacetaldehyde. In the cases of the dichloroethylenes and the trichloroethylene, it will be seen that the products can only result from a move of a chlorine atom. 66
Environmental Science & Technology
1,l-Dichloroethylene. In the case of the 1,l-dichloroethylene, a two-carbon reaction product, chloroacetyl chloride, is produced in large yield. A substantial amount of phosgene is also produced. In proportion to the amount of olefin reacted, the carbon monoxide and hydrogen chloride yields are substantially lower for 1,l-dichloroethylene than for ethylene or vinyl chloride. The ozone yield, however, is not lower. These observations are explainable in terms of the mechanisms outlined above. The phosgene obviously can be formed in the attack of ozone on the double bond Reactions 10 and 11 or in the equivalent uptake of an oxygen atom and an oxygen molecule. The chloroacetyl chloride is most likely formed in a rearrangement of a “hot” epoxide analogous to Reaction 15. In this case it must be concluded that the chlorine atom moves from one carbon to the other:
1,2-Dichloroethylene. In proceeding from 1,l-dichloroethylene to 1,2-dichloroethylene, the reaction products are most significantly changed. These changes provide an insight into the reaction mechanism, confirming the indicated sequence of reactions. First, it is seen that in this case with the starting olefin molecule having a chlorine atom on each carbon atom, no chloroacetyl chloride is formed. This confirms the belief that an intramolecular chlorine move takes place. In this case when the unstable epoxide undergoes a chlorine atom move, it produces dichloroacetaldehyde. This aldehyde will be further reacted, possibly producing more dichloroperoxyacetyl nitrate. The spectrum seems to indicate some small amount of a PAN-type compound by showing absorption a t the appropriate frequencies of 800 crnpi, 1300 cm-I and 1750 cm-I. Since the yield of chloroacetyl chloride is zero, and the yield of dichloroperoxyacetyl nitrate is a t best low, it follows that the carbon monoxide and hydrochloric acid yields should be high, as shown in Table I. These CO and HC1 yields are, in fact, two or three times greater than in the
cases of 1,l-dichloroethylene and trichloroethylene, both of which produce chloroacetylchlorides. Phosgene is not produced from 1,2-dichloroethylene, even though there is reason to believe that the chlorine atoms both can end up on the same carbon atom. This is explained by the sequence of reactions. The one-carbon carbonyl compounds are produced by the splitting of the olefin molecule. This and in competition with the molecular rearrangement, but not subsequent to the rearrangement. Trichloroethylene. Trichloroethylene as reactant yields products predicted by the general reaction mechanism discussed above. The observed HC1, CO, and phosgene can be ascribed to the attack of the various oxygenated species on the double bond. Since the starting material has a doubly chlorinated carbon atom, phosgene and an acyl chloride are reaction products, just as was the case with 1,l-dichloroethylene. In the trichloroethylene case, dichloroacetyl chloride is the product. Tetrachloroethylene. Tetrachloroethylene reacts much more slowly than the other olefins studied. The products produced fit the pattern predicted by the reaction mechanism. The chlorinated products are hydrogen chloride, phosgene, and trichloroacetyl chloride.
Literature Cited (1) Tariff Commission Reports, Miscellaneous Chemicals, Washington, D.C., 1974. ( 2 ) Molina, M. J., Rowland, F. S., Nature, 249,810 (1974). (3) Cicerone, R. J., Stolarski, R. S.,Walters, S., Science, 185, 1165 (1974). (4) Gay, B. W., Jr., Lonneman, W. A,, Bridboard, K., Moran, J . B., Ann N . Y Acad. Sci.. 246.286 (1975). (5) Jaeger, R. J., ibid.,’p 150. (6) Heicklen, J. P., Sanhueva. E.. Hisatsune. I. C.. Javantv. P. K. M., Simonaites, R., Hull, L. A., Blume, C. W., Mathias, E., “The
Oxidation of Halocarbons”, Final Report, EPA grant #R800949, May 1975. (7) Tuesdav. C. S..“The AtmosDheric Photooxidation of trans-2Butene and Nitric Oxide”, in ;‘Chemical Reactions in the Lower and Upper Atmosphere”, pp 1 to 49, Interscience, New York, N.Y., 1961. (8)Hanst, P. L., “Advances in Environmental Science and Technology”, Pitts and Metcalf, Eds., Published by J. Wiley and Sons, Inc., 1971. (9) Hanst, P. L., Lefohn, A. S., Gay, B. W., Jr., Appl. Spectros., 27.188 (1973). (10) ’Korth, M: W., Rose, A. H., Stahman, R. C., J . Air Pollut. Control Assoc., 14,168 (1964). (11) Wilson, D., Kopczynski, S. L., ibid., 18,160 (1968). (12) Saltzman, B. E., Anal Chem., 26,1949 (1954). (13) Byers, D. II., Saltzman, B. E., J . Am. Indust. Hyg. Assoc., 19, 261 (1958). (14) Altshuller, A. P., Miller, D. L., Sleva, S. F., Anal. Chem., 33, 621 (1961). (15) Hull. L. A.. Hisatsune. I. C.. Heicklen. 3.. Can. J . Chem.. 51. 1504 (1973). (16) Garvin. D.. Hamman. R. F.. “Chemical Kinetics Data Survey VII. Tables of Rate and Photochemical Data for Modelling of the Stratosphere” (Revised), Nat. Bur. Stand., 74-430, Washington, D.C., 1974. (17) Williamson, D. G., CvetanoviE, R. J., J . Am. Chem. SOC.,90, 3668 (1968). (18) Bufalini, J. J., Altshuller, A. P., Can. J . Chem., 43, 2243 (1965). (19) Leighton, P. A,, “Photochemistry of Air Pollution”, 300 pp, Academic Press, New York, N.Y., 1961. (20) Gay, B. W., Jr., Bufalini, J. J., Enuiron. Sci. Technol., 5 , 423 (1971). (21) Hisatsune, I. S., Heicklen, J., Can. J . Spectros., 18,77 (1973). (22) David McNelis, Ph.D. Thesis, University of North Carolina a t Chapel Hill, N.C., 1974. (23) Van Duuren, B. L., Ann. N . Y. Acad. Sci., 246, 258 (1975). (24) Gay, B. W., Jr. Noonan, R. C., Bufalini, J. J., Hanst, P. L., Enuiron. Sci. Technol. 10,38 (1975). Received for review December 23, 1974. Accepted August 28, 1975.
Analysis of Chromium in Natural Waters by Gas Chromatography R. J. Lovett‘ and G. Fred L e e * Institute for Environmental Sciences, University of Texas at Dallas, Richardson, Tex. 75080
An analytical method for the determination of chromium in natural waters by gas chromatography techniques has been developed. The chromium is chelated with HTFA and the chelate is extracted into benzene. The extract is injected into a gas chromatograph using an electron capture detector. A detection limit of 0.1 pgll. was found for chromium in natural water samples.
An accurate method for the determination of chromium in natural waters is essential if the true behavior of the element and its low level physiological effects are to be studied. The usual methods of aqueous chromium analysis, atomic absorption and uv-visible spectrophotometry are not generally feasible a t extremely low chromium levels. Even with sophisticated modifications, the most optimistic limit of detection of total chromium is 0.5 pg/l. using X-ray fluorescence following preconcentration on a chelating
’ Arizona State University, Tempe, Ariz.
resin (1).A flameless atomic absorption method can detect one pg/l. ( 2 ) .Chromium levels in natural waters are sometimes in the order of 0.05 pg/l. ( 3 ) .Therefore, routine analytical methods compatible with these levels should be developed. Much work has been done on the determination of chromium present a t low concentrations in other matrices by gas chromatography. Lunar samples ( 4 ) , ferrous alloys (5), and biological samples (6-9) have been analyzed for chromium by this method. The procedure involves chelation of the chromium with a P-diketone, l,l,l-trifluoro-2,4-pentanedione (also known as trifluoroacetylacetone and, hereafter, HFTA). Chromium TFA chelates have been found to be volatile enough for gas chromatographic analysis. By use of an electron capture detector, as little as 3 X lo-’* g of chromium can be detected (6). Based on a 5 - ~ injection, 1 the relative detection limit would be 0.01 pg/l. of chromium, which is quite sufficient for the analysis of a natural water. The problem of transferring chromium from a natural material into a benzene solution as a chelate can be accomVolume 10, Number 1, January 1976
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